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UNIVERSITY  OF  CALIFORNIA 


DEPARTMENT  OF  EDUCATION 


GIFT   OF  THE   PUBLISHER 

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No '         *  Received 


LIBRARY 

OF  THE 

UNIVERSITY  OF  CALIFORNIA. 


GIFT    OF 


Class 


An  Elementary 
Experimental  Chemistry 


BY 


JOHN  BERNARD  EKELEY,  A.M. 

* » 

Science  Master  at  St.  Paul's  School,  Garden  City 


SILVER,    BURDETT   &    COMPANY 
NEW  YORK.        BOSTON        CHICAGO 


COPYRIGHT,  1900,  BY 
SILVER.  BURDETT  &   COMPANY 


PREFACE 


THERE  are  two  things  that  the  study  of  a  science  should 
accomplish  for  the  student :  first,  the  development  of  the 
powers  of  observation ;  second,  a  knowledge  of  the  relations 
of  the  principles  and  facts  that  underlie  that  science.  The 
first  may  be  obtained  to  a  greater  or  less  degree  from  the 
study  of  any  science  by  the  experimental  method  ;  the  lat- 
ter however  may  often  be  only  imperfectly  acquired  even 
by  those  who  are  able  to  make  accurate  observations,  but 
who  fail  to  draw  the  correct  or  logical  conclusions.  It  is  the 
aim  of  this  book  to  aid  the  student  in  accomplishing  both 
these  things.  The  author  has  found  by  experience  (as,  in- 
deed, what  teacher  has  not  ?)  that  all  beginners  in  the  study 
of  chemistry  are  by  no  means  natural  adepts  at  making  the 
necessary  observations,  but  that  by  a  sufficiently  prolonged 
course  of  experiments  their  powers  in  this  direction  are 
easily  augmented.  The  difficulty  comes,  however,  in  ena- 
bling the, student,  after  the  observations  have  been  made,  to 
form  the  correct  conclusions,  and,  finally,  when  both  these 
things  have  been  done  throughout  the  subject,  to  round  it 
off  into  a  symmetrical  whole. 

To  accomplish  these  things,  the  author  believes  that  it  is 
necessary  to  help  the  student  considerably  by  emphasizing 
what  must  especially  be  marked  in  the  experiment,  and  to 
outline  the  course  of  thought  that  must  be  followed,  in  order 
that  the  student  may  feel  satisfied  at  the  close  of  the  experi- 


235384 


IV  PREFACE 

ment  that  he  has  made  some  definite  progress  at  least  in  the 
process  of  building  up  his  knowledge  of  the  subject. 

In  Parts  I.  and  II.,  this  method  is  followed.  Part  I.  is 
made  up  entirely  of  experiments  dealing  with  The  Prepara- 
tion and  Properties  of  the  Principal  Elements  and  Compounds. 
No  teaching  of  chemical  theory  whatever  is  attempted  in 
this  part.  Qualitative  equations  are  given,  showing  merely 
what  kinds  of  matter  have  been  concerned  in  each  chemi- 
cal change.  Words,  and  not  chemical  symbols,  are  used,  the 
conventional  chemical  equation  being  left  for  Part  II.  It 
is  hoped  that  the  student,  when  he  has  performed  the  ex- 
periments of  Part  I.,  may  have  a  good  idea  of  the  qualitative 
composition  of  the  principal  compounds,  and  of  how  he  may 
distinguish  them  from  one  another.  Each  experiment  gives 
him  a  knowledge  of  one  or  more  elements  or  compounds  ;  in 
doing  this,  no  compound  is  used  that  has  not  been  studied  in 
some  previous  experiment.  Thus  the  student's  knowledge  is 
built  up  step  by  step  by  the  inductive  method,  and  in  a  logi- 
cal manner,  until  he  has  made,  and  studied  the  properties  of 
those  elements  and  compounds  of  inorganic  chemistry  with 
which  he  is  most  concerned. 

The  author  believes  that  the  study  of  chemical  theory  is 
most  successfully  carried  on  after  the  student  has  prepared 
the  elements  and  compounds  and  has  studied  their  proper- 
ties. Hence  Part  II.  is  concerned  entirely  with  The  Laws 
and  Theories  of  Chemistry.  Special  emphasis  is  laid  upon  the 
difference  between  laws  and  theories.  The  laws  are  illus- 
trated by  experiments;  and  the  reasoning  upon  which  we 
base  our  belief  in  the  respective  explanatory  theories,  which 
most  students  find  so  difficult,  is  given  in  full.  This  gives 
an  accurate  idea  of  theoretical  chemistry  instead  of  the  hazy 
conceptions  that  are  so  often  obtained  when  the  theory  is 
scattered  throughout  a  series  of  merely  descriptive  experi- 


PREFACE  V 

ments.  The  use  of  chemical  symbols  and  the  writing  of 
reactions  are  now  taken  up  in  detail,  and  numerous  exam- 
ples are  given  to  illustrate  the  stoichiometrical  relations  of  the 
elements.  The  plan  of  deferring  the  writing  of  reactions 
until  late  in  the  course  is  in  line  with  the  most  recent  sug- 
gestions of  the  best  teachers  of  the  science.  Throughout 
the  book  the  revised  spelling  of  chemical  names  is  used. 
The  illustrations  of  apparatus  in  Parts  I.  and  II.  will  save 
the  teacher  valuable  time  that  would  otherwise  be  used  in 
answering  questions  on  manipulation. 

The  author  wishes  to  acknowledge  his  indebtedness  to 
Prof.  J.  F.  McGregory  of  Colgate  University,  and  to  Dr. 
Albert  C.  Hale  of  the  Boys'  High  School,  Brooklyn,  for 
many  valuable  suggestions,  and  to  his  colleague  at  St.  Paul's, 
Mr.  Arthur  De  L.  Ayrault,  for  his  painstaking  criticism  of 
the  English. 

JOHN    B.    EKELEY. 


INTRODUCTION 


TO  THE  TEACHER 

To  use  this  book,  especially  Parts  I.  and  II.,  to  the  best 
advantage,  the  order  of  the  experiments,  in  most  cases, 
should  not  be  changed.  Great  care  has  been  taken  to 
arrange  them  so  that  they  follow  one  another  in  logical 
order.  If  there  is  a  lack  of  time  to  do  them  all,  some  (the 
teacher  will  easily  see  what  ones)  can  be  omitted  without 
destroying  the  continuity  of  the  whole. 

Students  should  be  examined  by  means  of  recitations  or 
individual  examinations  after  each  ten  experiments  of  Part 
I.  At  St.  Paul's,  each  student  meets  the  instructor  for  a 
"  quiz  "  as  soon  as  ten  consecutive  experiments  have  been 
finished.  With  Part  II.,  however,  it  seems  best  to  supple- 
ment the  experimental  work  with  frequent  recitations. 

TO  THE   STUDENT 

Procure  a  blank  book  of  about  one  hundred  and  fifty  pages, 
with  durable  covers  (leather  back  and  corners).  Begin  your 
notes  on  the  sixth  page,  leaving  the  first  five  pages  blank 
for  a  future  "  table  of  contents."  On  the  left-hand  page, 
record  with  a  hard  pencil  the  notes  on  your  observations  as 
taken  at  the  time  of  the  experiment.  Do  not  fall  into  the 
habit  of  writing  your  original  notes  on  slips  of  paper,  but 
record  them  immediately  in  your  note  book.  When  your 


Vlll  INTRODUCTION 

original  notes  have  been  criticised  by  your  instructor,  study 
them  carefully.  After  you  are  satisfied  that  you  thoroughly 
understand  the  experiment,  write  out  in  ink  on  the  right- 
hand  page  in  your  best  English  a  detailed  description  of  the 
experiment  and  the  conclusions  that  you  have  reached. 

Before  attempting  an  experiment,  study  carefully  the  de- 
scription as  given  in  the  text,  so  that  you  are  sure  you  un- 
derstand what  you  are  going  to  do,  and  what  the  object  of 
the  experiment  is. 

Keep  your  laboratory  desk  and  apparatus  clean.  You 
will  often  be  tempted  not  to  do  this,  but  you  will  get  the 
best  results  from  your  work,  and  get  the  most  enjoyment 
from  it,  if  you  avoid  slovenly  habits  of  experimentation. 


CONTENTS 


PAGE 

Introduction , .          ...  vii 

PART    I. 

Preparation  and  Properties  of  Elements  and  Compounds    .     .     .  3-85 

Exp.     i.    Three  Conditions  of  Matter ...  3 

"       2.    Physical  and  Chemical  Changes 3 

"       3.    Mechanical  Mixture  and  Chemical  Compound  ....  4 

"       4.    Copper  and  Copper  Oxid ...  5 

"       5.    Mercury 6 

"       6.    Oxygen 7 

"       7.    Phosphorus 10 

8.    Carbon 12 

"       9.    Sulfur 13 

"     10.    Sodium 17 

"     n.    Potassium 18 

"     12.    Zinc ....          19 

"     13.    Magnesium 20 

"     14.    Iron .     .          .     .               ...  20 

"     15.    Hydrogen 21 

"     1 6.    Hydrogen  and  Oxygen 25 

"     17.    Zinc  and  Magnesium  Oxids,  and  Sulfuric  Acid  .     ...  27 

"     1 8.    Neutralization 28 

"     19.    Carbonic  Acid  and    Carbon   Dioxid  with   Sodium  and 

Potassium  Hydroxid 30 

Calcium 30 

Analysis  of  Marble 32 

22.  Chlorin ...  34 

23.  Hydrochloric  Acid  from  Sodium  Chlorid  and  Sulfuric 

Acid 36 


X  CONTENTS 

PAGE 

Exp.  24.    Analysis  of  Hydrochloric  Acid 37 

"  25.  Hydrochloric   Acid  with    Metals,  Hydroxids,  and    Car- 
bonates    38 

"  26.  Carbon  Dioxid  from  Marble  and  Hydrochloric  Acid  .     .  39 

"  27.    Preparation  of  Chlorin 40 

"  28.    Bromin 41 

"  29.    lodin 43 

"  30.    Potassium  lodid 44 

"  31.  Calcium  Fluorid  and  Hydrogen  Fluorid    ......  45 

"  32.    Sulfids 47 

"  33-    Other  Compounds  of  Carbon 49 

"  34.    Nature  of  Flame 54 

"  35.    Hard  and  Soft  Water 56 

"  36.    Nitrogen 57 

"  37.    Nitric  Acid 59 

"  38.    Neutralization  of  Nitric  Acid 62 

"  39.    Nitric  Oxid 63 

"  40.    Elements  in  the  Nascent  State 65 

"  41.    Ammonia 66 

"  42.    Ammonium  Chlorid    .     .' 67 

"  43.    Ammonia  from  Ammonium  Chlorid 68 

"  44.    Ammonium  Amalgam 69 

"  45.  Neutralization  of  Acids  with  Ammonium  Hydroxid  .     .  69 

"  46.    Nitrous  Oxid 70 

"  47.    Analysis  of  Nitrous  Oxid 72 

"  48.    Arsenic 72 

"  49.    Antimony 74 

"  50.    Bismuth 75 

"  51.    Cadmium 76 

"  52.    Mercury 77 

"  53-    Lead 77 

"  54-    Tin 79 

"  55.    Aluminum 79 

"  56.    Iron * 80 

"  57.    Nickel 8 1 

"  58.    Barium 81 

"  59.    Strontium 82 

"  60.    Silver 83 

"  61.    Gold 84 

"  62.    Platinum 85 


CONTENTS  XI 
PART    II. 

PAGE 

Laws  and  Theories  of  Chemistry 89-156 

Exp.    i.    The  Sum  of  the  Weights  of  the  Factors  in  a  Chemi- 
cal Change  equals  the  Sum  of  the  Weights  of  the 

Products 90 

"       2.    Law  of  Definite  Proportions  by  Weight 92 

"       3.    Law  of  Multiple  Proportions 93 

"       4.    Quantitative  Analysis  of  Sodium  Chlorid 94 

"•       5.    Barometer 98 

"      6.    Law  of  Boyle , 99 

"       7.    Law  of  Charles 101 

"       8.    Weight  of  a  Liter  of  Air 104 

"       9.    Weight  of  a  Liter  of  Carbon  Dioxid     ......  105 

"     10.    Weight  of  a  Liter  of  Hydrogen 106 

"     1 1.    Density 107 

"     12.    Density  of  a  Liquid  in  the  form  of  Vapor 107 

"     13.    Chemical  Equivalence 110 

'•     14.    Electrical  Equivalents 112 

"     15.    Specific  Heat  of  Lead 114 

'•     16.    Specific  Heat  of  Iron 1 16 

'•     17.    Law  of  Definite  Proportions  by  Volume 119 

"     1 8.    Molecular  Weight  of  Potassium  Chlorate 123 

"     19.    Molecular  Weight  of  Potassium  Chlorid 124 

"     20.    Writing  of  Reactions 141 

"     21.    Heat  of  Chemical  Action    .     .          147 

"     22.    Heat  of  Neutralization 149 

"     23.    Heat  of  Solution  and  of  Hydration 150 

"     24.    Dissociation  (Gaseous) 152 

"     25.    Dissociation  In  Liquids 153 

"     26.    Burning  of  Organic  Matter ;  Dry  Distillation     ....  154 

"     27.    Alcohol 155 

"      28.    Saponification    .          156 


Xll  CONTENTS 

PART    III. 

PAGE 

History,   Occurrence  and  Industrial  Applications   of    the 

Principal  Elements  and  Compounds 159-223 

Qualitative  Analysis 227-239 

Appendix  I 240 

Appendix  II 241 


PART  I. 

PREPARATION   AND   PROPERTIES   OF 
ELEMENTS  AND  COMPOUNDS 


AN   ELEMENTARY 
EXPERIMENTAL   CHEMISTRY 


PART    I. 

PREPARATION  AND  PROPERTIES  OF 
ELEMENTS  AND  COMPOUNDS 


EXPERIMENT    i 
Three  Conditions  of  Matter 

TAKE  a  piece  of  ice.  Note  its  properties.  After  crush- 
ing, heat  in  a  100  cc.  flask.  When  all  is  melted,  note  its 
properties.  Heat  the  water  obtained  to  boiling.  Does  any- 
thing escape  from  the  mouth  of  the  flask  ?  If  so,  state  its 
properties.  What  are  the  three  conditions  of  matter  ? 

EXPERIMENT    2 
Physical  and  Chemical  Changes 

a.  Dissolve  a  little  sugar  in  water.     Evaporate  until  all 
the  water  is  gone.     What  remains  ? 

b.  Heat  a  little  sugar  in  a  porcelain  dish.     Examine  the 
residue.     Is  it  still  sugar  ? 

3 


.. XPERIMENTAL  CHEMISTRY 

c.  Heat  an  iron  wire  in  a  flame.      Examine  it  again  after 
it  is  cool.      Is  it  still  iron  ? 

d.  Have  ready  a  few  small  iron  nails.      Pour  10   cc.  of 
water  into  a   100  cc.  flask,  and  then  add  5   cc.  of  sulfuric 

acid.  In  diluting  sulfuric  acid, 
always  pour  the  acid  into  the 
water.  If  you  pour  the  water 
into  the  acid,  it  will  spatter. 
Dangerous.  Place  the-  nails  in 
the  diluted  acid;  and,  after  all 

action  has  ceased,  filter,  evaporate,  and  allow  to  crystallize. 

Is  the  resulting  product  iron  ? 

State   what  you   conceive   to   be   the   difference   between 

physical  and  chemical  change. 


EXPERIMENT    3 
Mechanical  Mixture  and  Chemical  Compound 

An  element  is  a  substance  that  has  nof  been  separated  into  two 

or  more  dissimilar  substances. 
Sulfur  and  iron  are  elements. 

a.  Mix  equal  portions  of  flowers  of  sulfur  and  fine  iron 
filings.      Can  you  distinguish  the  particles  of  each  element 
on  looking  at  the  mixture  through  a  strong  magnifying  glass  ? 
Pass  a  magnet  through  the  mixture,  and  tap  it  lightly  on 
some  solid  body.     Place  a  little  of  the  mixture  in  a  test  tube, 
and  add  a  little   carbon  disulfid.     After  shaking  well  for  a 
few  moments,  pour  off  the  liquid  and  examine  the  residue. 
Is  it  iron  ?     Evaporate  the  liquid.     What  remains  ? 

b.  Place    the   remainder   of  the   mixture   in   a   small  test 
tube,  and  hold  the  bottom  of  the  tube  in  the  flame   of  i 
Bunsen   burner.      After  the   action   has   ceased,   break  thj 


ELEMENTS    AND    COMPOUNDS 


5 


tube,  and  examine  the  contents  by  the  same  tests  as  before 
heating.     Is  it  iron  ?     Is  it  sulfur  ? 

State  what  you  conceive  to  be  the  difference  between  a 
mechanical  mixture  and  a  chemical  compound. 


EXPERIMENT    4 
Copper  and  Copper  Oxid 

a.  Examine  some  copper  in  the  form  of  wire  and  sheet. 
State  as  many  of  its  properties  as  you  can  ;  i.e.,  color,  hard- 
ness, luster,  weight,  tenacity,  fusibility,  volatility,  etc. 

b.  Place  about  5  grms.  of  fine  copper  filings  in  a  porce- 
lain crucible,  and  weigh  carefully  to  i  c.  grm.     Heat  over  a 
Bunsen   burner,   stirring 

occasionally  with  a  clean 
iron  rod.  After  it  is  cool, 
weigh  again. 

Let  us  explain  the  in- 
crease in  weight.  Since 
matter  is  the  only  thing 
that  has  weight,  some 
kind  of  matter  must  have 
been  added  to  the  con- 
tents of  the  crucible. 
It  has  not  been  added 

to  the  crucible  itself;  for,  by  trying,  you  would  find  that  the 
weight  remains  unchanged.  This  matter  could  not  have 
come  from  the  heat,  since  heat  is  not  material ;  so  it  must 
have  come  from  the  air.  This  something  chemists  call 
oxygen.  Examine  the  contents  of  the  crucible  ;  you  will 
see  that  it  is  no  longer  copper.  A  chemical  change  has 
taken  place,  and  a  new  substance  called  copper  oxid  has 


6  AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

been  formed.  It  is  thus  named  to  indicate  the  chemical 
union  of  copper  and  oxygen. 

When  elements  unite  chemically,  the  resulting  substance 
is  called  a  compound.  The  process  of  combining  two  or 
more  elements  chemically  is  called  synthesis.  We  may  ex- 
press what  has  taken  place  by  the  following : 

*»«+<*•— 

Understand  that,  when  the  names  of  elements  are  written 
one  under  another,  the  combination  signifies  that  the  ele- 
ments are  chemically  combined. 


EXPERIMENT    5 
Mercury 

a.  Properties.     Examine  a  small  quantity  of  the  element 
mercury.      State  as  many  of  its  properties  as   you  can  dis- 
tinguish.      In    heating 
mercury,  be  careful  not 
to  breathe   any   of  the 
vapor. 

By  heating  mercury 
a  long  time  in  a  flask, 
it  is  possible  to  obtain 
a  red  powder  that 

weighs  more  than  the  original  mercury.  This  red  powder 
is  called  mercury  oxid,  and  is  formed  by  the  union  of  mer- 
cury and  oxygen  just  as  copper  oxid  is  formed  by  the  union 
of  copper  and  oxygen. 

b.  Analysis  of  Mercury  Oxid.     Place  a  small  quantity  of 
mercury  oxid  in  a  hard  glass  tube  of  about  15  mm.  diameter 


ELEMENTS    AND    COMPOUNDS  7 

and  closed  at  one  end.  Fit  the  tube  with  a  one-holed 
stopper  through  which  extends  a  glass  exit  tube  of  about 
4  mm.  bore.  By  means  of  a  piece  of  rubber  tubing,  con- 
nect this  with  a  glass  delivery  tube.  Clamp  the  hard  glass 
tube  on  a  standard  so  that  it  is  in  a  slightly  inclined  posi- 
tion. Have  ready  a  pneumatic  trough  with  a  small  beaker 
full  of  water  inverted  on  the  shelf.  Now  heat  the  mercury 
oxid  gently.  Gradually  increase  the  heat,  and  allow  the 
bubbles  to  escape  under  the  inverted  beaker  of  water. 
What  forms  on  the  sides  of  the  tube  ?  If  mercury  oxid  is 
composed  of  mercury  and  oxygen,  what  is  the  substance  in 
the  inverted  beaker  ? 

This  process,  the  separating  of  a  compound  into  two  or 
more  elements,  is  called  analysis. 

Mercury 

_  =  Mercury  +  Oxygen 

Oxygen 


EXPERIMENT    6 
Oxygen 

a.  Properties.     Examine  the  substance  in  the  jar  made  in 
the  last  experiment.     Place  a  ground-glass  plate  under  the 
jar,  and  remove  it  from  the  water.     Has   the    gas   formed 
color,  odor, "  or  taste  ?     Light  a  pine  splinter,  blow  out  the 
flame  so  that  a  spark  remains,  and  plunge  the  splinter  into 
the  jar.     What  happens  ?     \Vhat  is  the  substance  present  in 
the   air,  that  makes  fuel   burn  ?     Why   is   it  that  a   fire   is 
smothered  when  air  is  kept  away  from  it  ? 

b.  Preparation    from   Manganese    Dioxid.        Perform    an 
experiment    similar    to    Experiment   5  b,    using     manganese 
dioxid  (composed  of   manganese   and    oxygen)    instead    of 
mercury  oxid.     Is  a  metal  left  as  in  experiment  5  b  ? 


8  AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

c.  Preparation  from  Potassium   Chlorate.     Perform  a  sim- 
ilar experiment,  using  potassium    chlorate   instead  of  mer- 
cury  oxid.     Heat  gently  at  first,  beginning  at  the  top  of  the 
chlorate,  and  be  careful  to  keep  the  hands  from  under  the 
tube  in  case  it  should  break ;  for  melted  potassium  chlorate 
makes   a  serious  wound.     After  the  action  has  ceased,  dis- 
solve the  remaining  substance  in  hot  water,  and  evaporate 
to  dryness.     Preserve  this  residue,  and,  when  you  have  per- 
formed Experiment  25  c,  compare. 

d.  Preparation  from  a  Mixture  of  Potassium  Chlorate  and 
Manganese  Dioxid.     Perform   a  similar  experiment,  using  8 
grms.   of   potassium    chlorate    and    2    grms.   of   manganese 
dioxid.     Mix    the  two  thoroughly.      Begin   heating   at  the 
top    of   the   mixture.     When   action    has    ceased,   add    hot 
water   and  transfer  to   a  beaker.      Be   sure  to  get  all   the 
residue.      Weigh   a   filter  paper    carefully.      Filter   the    hot 
liquid,  being  careful  to  throw  it  all  and  the  black  residue 
upon  the   filter  paper.     Wash  with  hot  water  from  a  wash 
bottle  five  or  six  times.     Dry  the  filter  in  an  oven  at  a  tem- 
perature of  about  100°.    Weigh  again.     Subtract  the  weight 
of  the  filter   paper  from  the  weight  of  the  filter  paper  and 
black  residue.     How  does  the  result  compare  with  the  ori- 
ginal weight  of  the  manganese  dioxid  ?     Has  the  manganese 
dioxid  undergone  any  change  chemically  ?     Evaporate  the 
filtrate  (the  liquid  that  ran  through  the  paper)  to  dryness  in 
a  weighed  porcelain  dish.     Weigh,  and  account  for  the  loss 
of  weight  of  the  chlorate. 

e.  Preparation    by  Electrolysis  of   Water.     Have   ready   a 
vessel   fitted  with  two  platinum  electrodes.     In£o  this  pour 
sufficient  water  acidified  with  sulfuric  acid  (i  part  acid  to 
20  of  water),*  so  that  it  will  somewhat  more  than  cover  the 
electrodes.     Fill  two  glass  tubes  graduated  in  cubic  centi- 

*  The  acid  is  only  added  to  aid  the  water  in  conducting  the  current. 


ELEMENTS  AND  COMPOUNDS  9 

meters  with  some  of  the  water,  and  invert  them  in  the  vessel 
so  that  the  mouth  of  each  tube  covers  an  electrode.  Place 
the  vessel  in  circuit  with  a  battery  of  two  bichromate  cells. 
Note  the  evolution  of  gas  at  each  electrode.  After  a  few 
minutes,  disconnect  the  cells,  and  measure  the  relative  vol- 
umes of  the  gases.  Test  the  lesser  one  with  a  glowing 


splinter.  What  is  it  ?  Apply  a  lighted  match  to  the  greater, 
and  note  the  result.  Here  we  have  a  new  gas,  which 
chemists  have  been  unable  to  decompose  into  anything  else. 
It  has  been  named  hydrogen.  We  now  have  experimental 
proof  as  to  the  composition  of  water. 


IO        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Remark.  If  a  silent  discharge  of  electricity  is  passed 
through  dry  oxygen,  the  gas  will  decrease  in  volume  one- 
third,  and  will  acquire  greater  chemical  activity.  Some 
change,  then,  has  occurred  in  the  gas  without  the  addition 
of  any  other  kind  of  matter.  When  elements  can  thus  be 
changed  by  any  means,  they  are  said  to  exist  in  allotropic 
forms.  This  second  form  of  oxygen  is  called  ozone.  It  has 
a  faint  peculiar  odor,  often  noticed  near  electric  machines 
when  operating,  or  just  after  a  discharge  of  lightning. 


EXPERIMENT  7 
Phosphorus 

Phosphorus  occurs  in  two  allotropic  forms,  the  yellow  and 
the  red.  The  yellow  must  be  stored  under  water,  as  it 
ignites  at  a  temperature  of  40°  in  the  air. 

a.  Properties.  Examine  both  red  and  yellow  phosphorus. 
Be  very  careful  in  handling  the  yellow  variety,  since  it  ignites 
so  easily,  and  a  burn  from  it  heals  with  difficulty.  Never 

handle  it  with  the  fingers, 
and  always  cut  it  under 
water.  State  as  many 
properties  as  you  can  of 
each  kind. 

b.  Preparation  of  Phos- 
phoric Oxid  by  Burning 
Phosphorus  in  air.  Have 
ready  a  quick-sealing  fruit 

jar.  The  rubber  washer  should  be  greased  with  vaseline. 
On  a  support  in  the  jar,  hang  a  deflagrating  spoon  con- 
taining a  piece  of  ignited*  phosphorus  about  the  size  of 
*  The  phosphorus  may  be  placed  in  the  jar,  and  then  ignited  by 
means  of  a  strong  lens. 


ELEMENTS    AND    COMPOUNDS  I  I 

a  pea.  Close  the  jar  air  tight.  Note  what  forms  in  the 
jar.  This  product  is  called  phosphoric  oxid.  Open  the  jar 
under  water.  Was  any  of  the  air  in  the  jar  used  up  ? 
Does  the  oxid  dissolve  in  water  ? 

Phosphorus  +  Oxygen  - 


Remark.  Another  oxid  of  phosphorus  exists,  called  phos- 
phorous oxid.  It  is  formed  when  phosphorus  is  burned  in  an 
insufficient  supply  of  air.  It  is  a  white  powder  with  an 
odor  resembling  garlic.  When  heated  in  the  air,  it  becomes 
the  oxid  you  have  just  made  ;  this  shows  that  it  differs  from 
phosphoric  oxid,  in  that  it  contains  less  oxygen. 

c.  Preparation  of  Phosphoric  Oxid  by  burning  Phosphorus  in 
Oxygen.    Perform  a  similar  experiment,  using  a  jar  of  oxygen 
instead  of  air.     In  this  case,  it  is  better  to  ignite  the  phos- 
phorus in  the   jar  by  means  of  a  burning  glass.     How  does 
the  product  compare  with  that  formed  in  b  ? 

d.  Phosphoric  Acid.     Dissolve  the  oxid  prepared  in  c  in  a 
small  amount  of  water.     Taste  a  drop  of  the  solution.     Also 
note   its   effect  upon  blue    litmus    paper.     This   product  is 
called  phosphoric  acid  ;  and,  since  it  is  made  of  phosphoric 
oxid  and  water,  it  can  contain  only  the  elements  hydrogen, 
phosphorus,  and  oxygen. 


Hydrogen       Phosphorus 
*  +  '  =  Phosphorus 

Oxygen  Oxygen 


Remark.  When  phosphorus  is  burned  in  an  insufficient 
supply  of  moist  air,  another  acid  of  phosphorus  is  formed 
which  has  a  faint  odor  like  garlic.  This  is  called  phospho- 
rous acid.  It  differs  from  phosphoric  acid  in  that  it  contains 
less  oxygen.  Note  the  signification  of  the  endings  ic  and 
ous. 


12         AN     ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Definition  of  an  Acid.  For  the  present  it  will  be  sufficient 
to  define  an  acid  as  a  compound  of  hydrogen,  or  else  of 
hydrogen  and  oxygen,  with  a  non-metallic  element,  which 
compound  has  usually  a  sour  taste  and  turns  blue  litmus 
paper  red. 

EXPERIMENT  8 
Carbon 

a.  Properties.     Examine,  and  state  the  properties  of,  the 
following  allotropic  forms  of  carbon:   charcoal,  gas  carbon, 
bone  black,  graphite,  and  soot. 

b.  Carbon  Dioxid.     Burn  a  piece  of  charcoal   in  a  defla- 
grating spoon  in  a  jar  of  oxygen,  as  in  Experiment  7  b.  *  The 
charcoal  must  be  well  ignited  before  being  placed  in  the  jar. 
Open  the   jar  under  water.     Does  the  cover  stick  ?     What 
can  you  say  of  the  volume  of  the  resulting  carbonic  oxid 
compared  with  the  original  volume  of   oxygen  ?     Close  the 
jar,   allowing  a  little  water  to   enter,  and  shake  vigorously. 
Open  again  under  water.    Is  carbonic  oxid  soluble  in  water  ? 
In  the  same  way,  prepare  another  jar  of  the  gas.     Plunge  a 
lighted  splinter  into  the  gas.     Does  the  gas  burn?     Does 
the  splinter  continue  to  burn  ?     The   name  carbon  dioxid  is 
given  to  this  gas  for  reasons  that  will  appear  later. 

-  Carbon 

Carbon  +  Oxygen  =  Qxygen 

c.  Carbonic  Acid.     Taste  the   solution   formed  in  b,  and 
test  it  with  blue  litmus  paper.     Here  we  have  a  second  acid, 
which  has  been  named  carbonic  acid.      Transfer  to  a  test  tube 
and  heat  to  boiling.     Taste,  and  test  with  litmus  paper  again. 
Is  carbonic  acid  a  stable  compound,  i.  e.,  is  it   hard  to  de- 
compose ?     What  has  become  of  the  carbon  dioxid  ? 


ELEMENTS  AND  COMPOUNDS  13 


Carbon        Hydrogen  = 
Oxygen         Oxygen 

Remark.  In  Experiment  7  b,  and  8  £,  we  have  observed 
the  phenomenon  of  combustion.  Combustion  is  the  chemical 
union  of  substances,  accompanied  by  the  evolution  of  light 
and  heat.  The  substances  that  unite  may  be  any  substances 
whatever,  but  ordinarily  we  apply  the  term  to  the  union  of 
substances  with  the  oxygen  of  the  air.  The  substance  that 
unites  with  the  oxygen  is  called  the  combustible,  while  the 
oxygen  is  said  to  support  combustion.  The  ease  with  which 
elements  unite  with  oxygen  varies.  Some,  such  as  phospho- 
rus, require  only  a  slight  rise  in  temperature,  wrhile  others, 
like  carbon,  require  a  comparatively  high  one.  The  temper- 
ature at  which  a  substance  takes  fire  in  air  is  called  its 
kindling  temperature.  This  temperature  is  constant  for  any 
particular  substance. 


EXPERIMENT  9 

Sulfur 

a.  Properties.     After  examining,  state  as  many  properties 
as  you  can  of  roll   sulfur.     Place  about   10  grms.  in  a  test 
tube,  and  heat  over  a  Bunsen  burner,  heating  gently  at  first 
and  gradually  increasing  the  temperature.     Note   carefully 
the  changes  that  the  sulfur  passes  through,  especially  in  re- 
gard to  its  color  and   consistency.     Note  also  the  ease  with 
which  sulfur  takes  fire. 

b.  Allotropic  Forms  of  Sulfur. 

i.    In  a  beaker  put  about  60  grms.  of  roll  sulfur,  and  heat 
gently  until  all  is  melted.     When  all  is  melted,  allow  it  to 


14          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

cool.  As  soon  as  crystals  have  just  covered  the  surface  of 
the  liquid,  break  a  hole  in  the  crust,  and  quickly  pour  out 
the  melted  sulfur.  Note  the  shape  and  color  of  the  crystals 
that  remain  in  the  beaker.  Allow  it  to  stand  for  a  few  days, 
and  examine  again. 

2.  Pulverize  a  small  lump  of  roll  sulfur,  and  shake  it  up 
with   carbon   bisulfid   in   a  test  tube.     Keep   away  from  a 
flame,  since  carbon  bisulfid  is  very  inflammable.     When  it 
is  dissolved,  pour  it  into  an  evaporating  dish  and  allow  it  to 
crystallize.    Note  the  color  and  shape  of  the  crystals  formed. 

3.  In  a  test  tube,  heat  to  boiling  about  15   grms.  of  roll 
sulfur,  and  quickly  pour  the  liquid  into   a  beaker  of  cold 
water.     State   the   properties  of  the   product.     Allow   it  to 
stand  for  a  few  days,  and  examine  again. 

c.  Sulfur  Dioxid.     Burn   a  small   amount  of  sulfur  in  a 
deflagrating  spoon  in  a  jar  of  oxygen.     Note  the  color  of  the 
flame.     After   action   ceases,  open   under  water.     Does  the 
cover  stick?     What  can  you  say  of  the  volume  of  the  gas 
compared  with  the  original  volume  of  oxygen  ?     Get  its  odor 
and  color,  if  any.     Allow  a  little  water  to  enter  the  jar,  seal 
again,  and  shake.    Open  again  under  water.     Does  the  cover 
stick  ?     Is  the  gas  soluble  in  water  ?     We  shall  call  this  gas 
sulfur  dioxid  for  reasons  that  will  appear  later. 

Sulfur 
Sulfur  +  Oxygen  =  _ 

Oxygen 

d.  Sulfurous  Acid.     Taste    the  solution  made  in  b,  and 
test  it  with  blue  litmus  paper.     This  is  called  sulfurous  acid. 


Sulfur 
Oxygen        Oxygen 


e.    Sulfur  Trioxid.     Have  ready  a  gas  holder  full  of  oxy- 
gen.     Apply  to  the  instructor  for  an  apparatus  for  obtaining 


ELEMENTS  AND  COMPOUNDS  15 

a  steady  stream  of  sulphur  dioxid.  (See  g.)  By  means  of 
rubber  tubing  and  a  glass  Y  tube,  allow  the  oxygen  and 
sulfur  dioxid  to  come  together.  In  order  to  dry  the  mixed 
gases,  pass  them  by  means  of  rubber  tubing  through  a  catch 
bottle  containing  concentrated  sulfuric  acid,  which  has  the 
property  of  extracting  moisture  from  gases  as  they  bubble 
through  it.  Then  pass  the  dry  gases  through  a  hard  glass 
tube,  about  15  cm.  long  and  i  cm.  bore,  containing  platinized 


asbestos.  Do  not  pack  the  asbestos  in  the  tube,  but  place 
it  in  loosely.  Have  the  tube  containing  the  asbestos  at- 
tached to  a  test  tube  fitted  with  a  two-holed  stopper  contain- 
ing an  entrance  and  an  exit  tube.  Let  the  entrance  tube 
extend  down  to  within  two  centimeters  of  the  bottom.  Pack 
the  test  tube  in  a  mixture  of  ice  and  salt.  Now  heat  the 
asbestos  to  redness.  The  red-hot  platinized  asbestos  simply 
aids  in  the  chemical  union  of  the  two  gases.  After  three  or 
four  minutes,  remove  the  test  tube  from  the  ice.  What  is 


1 6         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

the  nature  of  the  crystals  formed  ?  Note  especially  the 
effect  of  warmth  upon  them.  This  product  is  called  sulfur 
trioxid.  Since  it  is  formed  by  the  chemical  union  of  sulfur 
dioxid  and  oxygen,  it  differs  from  sulfur  dioxid  in  that  it 
contains  more  oxygen. 

Sulfur  Sulfur 

Oxygen  +°Xygen  =  Oxygen 

/  Sulfuric  Acid.  Allow  a  few  drops  of  water  to  fall  upon 
the  oxid  made  in  e.  Get  the  properties  of  the  resulting 
compound,  especially  its  taste  (dilute 
a  little  with  considerable  water  be- 
fore tasting),  and  action  upon  litmus 
paper.  This  compound  is  called  sul- 
furic acid,  and  evidently  differs  from 
sulfurous  acid  in  that  it  contains 
more  oxygen.  Examine  a  little  com- 
mercial sulfuric  acid. 


Oxygen      Oxygen 

Oxygen 

g.  Su  Ifu  r 
Dioxid  from 
Sulfuric  Acid 
and  Copper. 
In  a  250  cc. 
]  flask,  place 
about5ogrms. 
of  sheet  cop- 
per clippings. 

Fit  the  flask  with  a  two-holed  stopper  containing  an  exit 
tube  and  a  funnel  tube  that  extends  down  to  the  copper. 
Pour  concentrated  sulfuric  acid  into  the  flask  till  it  somewhat 


ELEMENTS  AND  COMPOUNDS  I/ 

more  than  covers  the  end  of  the  funnel  tube.  Heat  with  a 
Bunsen  burner,  and  collect  a  jar  full  of  the  gas  by  means  of 
a  glass  tube  passing  to  the  bottom  of  the  jar.  Note  that  it 
is  the  same  gas  as  you  obtained  on  burning  sulfur  in  oxygen, 
i.e.,  sulfur  dioxid. 


EXPERIMENT    10 
Sodium 

The  element  sodium  must  be  stored  under  naphtha,  to  pre- 
serve it.  When  you  wish  to  use  any,  apply  to  the  instructor 
for  it.  Never  handle  it  with  the  fingers.  Use  forceps. 

a.  Properties.     Examine   a   piece  of  clean   sodium.     Get 
its  properties,  —  hardness,  color,  etc.     Is  it  a  metal  ?     Make 
a  fresh  cut  on  it  with  a  knife,  and  note  the  appearance  of 
the  clean  surface  for  a  few  moments.     Throw  a  small  piece 
upon  water,  and  note  the  result. 

b.  Sodium  Oxid.     In  a  crucible,  heat  a  piece  of  clean  so- 
dium about  the  size  of  a  pea.     As  it  burns,  note  the  color  of 
the  flame.     Get  the  properties  of  the  resulting  sodium  oxid. 

Sodium 
Sodium  +  Oxygen  = 


c.  Sodium  Hydroxid.  When  the  oxid  made  in  b  is  cooled, 
allow  a  few  drops  of  water  to  fall  on  it.  Note  the  sputtering 
sound.  Put  a  drop  of  the  solution  on  your  finger,  and  note 
the  greasy  feeling.  Test  the  liquid  with  red  litmus  paper. 
Evaporate  it  to  dryness,  and  allow  it  to  stand  for  some  time, 
after  which  note  that  moisture  has  collected  upon  the  sub- 
stance. 

Since  sodium  oxid  has  united  with  water,  the  resulting 
compound  must  be  composed  of  sodium,  oxygen,  and  hydro- 


1  8         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

gen.  It  is  called  sodium  hydroxid  or  sodium  hydrate.  This 
is  the  first  of  a  class  of  compounds  called  bases,  which  you 
will  study. 

Sodium 
Sodium      Hydrogen 


Definition  of  a  Base.  For  the  present  it  is  sufficient  to  de- 
fine a  base  as  a  compound  of  a  metallic  element  with  oxy- 
gen, or  with  oxygen  and  hydrogen.  Bases  that  turn  red 
litmus  paper  blue  are  called  a/kalis. 

d.  Sodium  Amalgam.  Weigh  out  20  grms.  of  mercury, 
and  place  it  in  an  iron  pan  under  a  hood.  Heat  it  to  about 
200°.  The  temperature  may  easily  be  observed  by  means  of 
a  thermometer.  Take  2  grms.  of  clean  sodium,  and  by 
means  of  a  pair  of  long  forceps  drop  it  into  the  mercury. 
Step  back  instantly,  since  mercury  and  sodium  unite  with 
great  violence,  and  poisonous  vapors  of  mercury  are  evolved. 
The  product  is  called  sodium  amalgam. 


EXPERIMENT    n 
Potassium 

Potassium  is  stored  in  a  way  similar  to  sodium,  and  must 
be  handled  with  equal  care. 

a.  Properties.     Note  the  properties  of  potassium  as  you 
did  those  of  sodium.     Compare  the  metal  with  sodium. 

b.  Potassium  Oxid.     Perform  an  experiment  with  a  piece 
of  potassium,  corresponding  to  Experiment  10  b.     Note  the 
color  of  the  flame. 

Potassium 
Potassium  +  Oxygen  =  _ 

Oxygen 


ELEMENTS  AND  COMPOUNDS  IQ 

c.  Potassium  Hydroxid.  Perform  an  experiment  with 
potassium  corresponding  to  Experiment  10  c.  How  does 
the  compound  here  made  differ  in  chemical  composition 
from  that  made  in  Experiment  10  ^?  Here  we  have  another 
compound  belonging  to  the  class  of  bases.  It  is  called 

potassium  hydroxid. 

Potassium 
Potassium  .  Hydrogen       _ 

Oxygen      +  Oxygen     =  <*?•* 
Hydrogen 

EXPERIMENT    12 
Zinc 

a.  Properties.      Examine,  and  note  the  properties  *  of,  the 
metal  zinc  in  its  various  forms,  —  sheet,  stick,  granular,  and 
dust. 

b.  Zinc  Oxid.     Heat  a   small  piece  of  zinc   in  a  crucible, 
stirring  with  an   iron   rod.     Note   the  color  of  the  flame  as 
the  zinc  burns,  the   color  of  the   oxid  when  hot  and  when 
cold,  and  the  peculiar  woolly  appearance  that  it  assumes. 

Zinc  +  Oxygen  =  _ 

Oxygen 

c.  Zinc  Oxid  and   Water.     Try  the  effect  of  water  on  a 
little  oxid  of  zinc,  and  test  it  with  litmus  paper  turned  red 
with  carbonic  acid.     Filter,  and  evaporate  the  filtrate.     Is 
zinc  oxid   soluble   in  water  ?     Does  zinc  oxid  form  an  acid 
or  a  base  ? 

*  Zinc  is  peculiar  in  that  its  physical  properties  vary  considerably 
with  heat.  At  the  ordinary  temperature,  zinc  is  rather  brittle,  but,  if 
heated  to  100°,  it  becomes  malleable  and  can  be  rolled  into  sheets.  At 
205°  it  becomes  so  brittle  that  it  can  be  powdered.  It  melts  at  400°, 
and  boils  at  about  1000°.  Granular  zinc  is  formed  when  the  molten 
metal  is  allowed  to  fall  drop  by  drop  into  cold  water.  Zinc  dust  is  ob- 
tained when  the  vapors  of  boiling  zinc  are  suddenly  condensed  in  the 
absence  of  air. 


2O        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


EXPERIMENT    13 
Magnesium 

a.  Properties.     Examine  the  element  magnesium  both  in 
the  form  of  ribbon  and  powder  (flash  light  powder).     State 
its  properties. 

b.  Magnesium  Oxid.     Oxidize  a  little  magnesium  in  a  cru- 
cible, noting  the  color  of  the  oxid  when  hot  and  when  cold. 
How  could  you  distinguish  magnesium  oxid  from  zinc  oxid  ? 

Magnesium 
Magnesium  +  Oxygen  =  Oxygen 

c.  Magnesium  Oxid  and  Water.  Treat  magnesium  oxid  with 
water  as  you  did  zinc  oxid  in  Experiment  12  c.     Does  mag- 
nesium oxid  form  an  acid,  or  a  base  ? 


EXPERIMENT    14 
Iron 

a.  Properties.     Examine,  and  state  the  properties  of,  iron 
both  in  the  form  of  nails  and  in  the  form  of  powder  (iron  by 
hydrogen). 

b.  Iron   Oxid.      (Burning  iron   in  air.)     Place   about    10 
grms.  of  "  iron  by  hydrogen  "  in  a  crucible,  and  heat,   stir- 
ring   occasionally.      The   resulting    compound    is    evidently 

iron  oxid. 

Iron 
Iron  -f  Oxygen  =  _ 

Oxygen 

c.  Iron   Oxid  and   Water.     Try  the   effect   of  water  upon 
iron  oxid. 

d.  Iron  Oxid.     (Burning  iron  in  oxygen.)     Fill  a  jar  with 


ELEMENTS  AND  COMPOUNDS 


21 


oxygen  gas.     Pierce  a  hole  through  a  flat  cork  large  enough 

to  cover  the  mouth  of  the  jar.     Through  this  hole,  insert  a 

watch  spring   from  which 

the  temper  has  been  taken 

by  heating  in  the  flame  of 

a  Bunsen  burner.    Around 

one  end  of  the  spring,  wind 

a  little  cotton  string,  and 

soak  this  in  melted  sulfur. 

Ignite     the     sulfur,     and 

plunge  it  into  the  jar  of 

oxygen,  making  the   cork 

cover   the    mouth   of    the 

jar.      As    the    spring    is 

consumed,  feed  it  through 

the      hole      until      action 

ceases.     What  compound 

is  formed  ? 

EXPERIMENT    15 
Hydrogen 

a.  Hydrogen  by  Electrolysis  of  Water.     Properties.     Make 
hydrogen  again  by  electrolysis  of  water,  as  in  Experiment  6  e. 
Note  its  properties. 

b.  Hydrogen  from  Steam  and  Hot  Iron.     Place  about  25  or 
30  grms.  of  dean  iron  filings  in  the  middle  of  a  piece  of  half- 
inch  gas  pipe  about  two  feet  long.     Be  careful  that  the  filings 
do  not  stop  up  the  pipe.     Fit  the  ends  with  one-holed  stop- 
pers containing  glass  tubes  about  5  mm.  bore.    To  one  end  of 
the  pipe  connect  a  500  cc.  flask  about  half  full  of  water,  and 
to  the  other  a  rubber  tube  fitted  with  a  glass  leading  tube. 
Heat  the  gas  pipe  red  hot  with  a  blast  lamp,  and  then  boil 


22        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

the  water  in  the  flask.  Collect  the  gas  that  issues  from  the 
leading  tube  in  test  tubes  over  water.  Compare  the  gas 
with  the  hydrogen  obtained  in  a.  Be  careful  to  have  no 
flame  near  the  leading  tube,  since  hydrogen  mixed  with  air 
is  an  explosive  compound. 

Let  us  see  what  has  taken  place.     We  know  from  Experi- 
ment 6  e  that  water  is  composed  of   oxygen  and  hydrogen, 


and  from  Experiment  \\b  we  know  that  iron  will  unite 
chemically  with  oxygen.  Examine  the  substance  left  in  the 
gas  pipe,  and  recognize  it  as  iron  oxid. 

Hydrogen  ,  Iron 

Oxygen     +  Ir0n  =  Oxygen  +  =**<*" 

c.  Hydrogen  from  Sodium  and  Water.  In  a  cage  made  of 
fine  wire  gauze  and  fitted  with  a  handle,  place  a  small  piece 
of  metallic  sodium  about  the  size  of  a  pea.  Have  ready  a 
pneumatic  trough  with  an  inverted  beaker  full  of  water  on 
the  shelf.  Hold  the  cage  containing  the  sodium  under 
water.  After  a  few  bubbles  have  escaped,  hold  it  under 


ELEMENTS  AND  COMPOUNDS  23 

the  beaker  and  collect  the  gas  evolved.  When  action  has 
ceased,  recognize  the  gas  as  hydrogen. 

Into  a  small  beaker  full  of  water,  throw  a  few  small  pieces 
of  sodium,  and  test  the  liquid  afterwards  with  red  litmus 
paper. 

Sodium    has    so   strong  an   attrac- 


tion for  oxygen  that  it  takes  oxygen 


out  of  water,  leaving  the  hydrogen. 

The  resulting  oxid  unites  with  water 

and    forms    the    hydroxid.      If  there 

is    an    insufficient    supply   of    water, 

the  sodium  bursts  into  the  characteristic  yellow  flame.     To 

show  this,  place   a  small   piece  of  sodium   on  a  wet  piece 

of  filter  paper,  and  note  the  result. 

,   Hydrogen      Sodium 
(i)  Sodium  +  =  0xygen  +  Hydrogen 


Sodium     Hydrogen        °dlu™ 
Oxygen      Oxygen 


d.    Hydrogen  from  Potassium  and  Water.     Perform  an  ex- 
periment similar  to  c,  using  potassium  instead  of  sodium. 

Hydrogen      Potassium 

(i)  Potassium  +    J  +  Hydrogen 

Oxygen         Oxygen 

Potassium 
Potassium      Hydrogen 

Oxygen          Oxygen 


e.  Hydrogen  from  Zinc  and  Sulfuric  Acid.  In  a  500  cc. 
flask  fitted  with  a  two-holed  stopper  containing  an  exit  and 
a  funnel  tube,  place  about  100  grms.  of  granular  zinc.  Add 
enough  sulfuric  acid  (i  part  acid  and  5  parts  water)  to  cover 
the  zinc.  Catch  the  gas  evolved  over  water,  and  recognize 
as  hydrogen.  Be  sure  to  have  no  flame  near  the  flask. 


24         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

After  no  more  zinc  will  dissolve,  filter  the  contents  of  the 
flask ;  then,  after  the  solution  has  evaporated  somewhat, 
allow  it  to  crystallize.  Note  the  properties  of  the  crystals. 
Place  a  dry  crystal  in  a  glass  tube  closed  at  one  end,  and 
heat  gently  in  the  flame  of  a  Bunsen  burner.  Notice  the 
water  that  collects  on  the  upper,  cool  part  of  the  tube. 
Does  the  crystal  remain  intact  ?  This  water,  then,  seems 
in  some  way  to  be  necessary  to  the  crystal,  and  is  called 
"  water  of  crystallization." 

Since  sulfuric  acid  is  composed  of  hydrogen,  sulfur,  and 
oxygen,  and  since  in  the  experiment  the  zinc  disappears  and 


mu^    ^     m 

fSii'i'l  ^Afri^Ki^iMii^0  •    :  -•"-'  I         Hw^" 

^^^^^^^^^^^PBBBBHBRl^ 


hydrogen  is  evolved,  it  is  evident  that  the  zinc  takes  the 
place  of  the  hydrogen  in  the  sulfuric  acid.  The  crystals  are 
therefore  composed  of  zinc,  sulfur,  and  oxygen,  together  with 
the  water,  which  holds  the  combined  elements  in  crystalline 
form.  When  this  water  is  driven  off,  of  course  only  zinc,  sul- 
fur, and  oxygen  are  left,  in  the  form  of  a  white  powder.  This 
is  called  zinc  sulfate. 

Hydrogen  Zinc 

Zinc  +  Sulfur       =  Hydrogen  +  Sulfur 
Oxygen  Oxygen 


ELEMENTS    AND    COMPOUNDS 


/  Hydrogen  from  Iron  and  Sulf uric  Acid.  Perform  a  simi- 
lar experiment,  using  iron  nails  instead  of  zinc.  What  is  the 
color  of  the  crystals  obtained  ?  Of  what  are  they  composed  ? 
What  would  you  naturally  name  them  ? 


•     Hydrogen 
Iron  +  Sulfur 
Oxygen 


Iron 

=  Hydrogen  +  Sulfur 
Oxygen 


EXPERIMENT    16 
Hydrogen  and  Oxygen 

a.    Synthesis  of  Water  by  Burning  Hydrogen  in  Air.     Set 
up  an  apparatus  as  in  Experiment  15  e.    To  dry  the  hydrogen, 


pass  it  through  a  catch  bottle  containing  concentrated  sulfu- 
ric  acid,  and  have  the  delivery  tube  drawn  to  an  opening  of 
about  i  mm.  diameter.  After  the  gas  has  been  escaping 


26         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

for  a  few  minutes,  collect  a  test-tubeful  over  water.  If,  on 
ignition,  no  explosion  is  heard,  it  is  safe  to  light  the  hydro- 
gen as  it  escapes  from  the  apparatus.  Never  light  hydrogen 
from  a  generator  without  taking  this  precaution.  Hold  a 
dry  bell  jar  over  the  flame,  and  keep  the  jar  cool  with  a  cloth 
wrung  out  in  cold  water.  Notice  the  water  forming  on  the 
inside  of  the  jar.  Here  we  have  further  proof  of  the  com- 
position of  water.  In  Experiment  6  <?,  we  proved  by  analy- 
sis that  water  is  composed  of  hydrogen  and  oxygen.  Here 
the  same  thing  is  proved  by  synthesis. 

Remark  i.  Water  is  also  formed  by  the  oxidation  in  the 
air  of  substances  containing  hydrogen.  Ignite  a  piece  of 
wood,  and  hold  the  flame  near  a  cold  glass  plate.  Note  the 
moisture  condensed  upon  the  plate. 

By  the  slow  oxidation  in  the  human  body  of  substances 
containing  hydrogen,  water  is  formed.  Breathe  upon  a  cold 
glass  plate,  and  note  the  moisture  condensed  upon  it.  For 
a  second  product  of  oxidation  in  the  human  body,  see  Ex- 
periment 20  d. 

Remark  2.  If  the  earth  were  covered  with  an  atmosphere 
of  hydrogen  instead  of  oxygen,  a  stream  of  oxygen  issuing 
from  a  jet  would  burn,  if  ignited,  just  as  the  stream  of  hy- 
drogen burns  in  this  experiment. 

b.  Reduction  of  Copper  Oxid  by  Hydrogen.  In  a  combus- 
tion tube,  place  about  5  grms.  of  copper  oxid.  Have  ready  a 
hydrogen  generator  connected  with  a  catch  bottle  containing 
concentrated  sulfuric  acid,  and  attach  this  to  one  end  of  the 
combustion  tube.  To  the  other  end,  attach  a  dry  test  tube 
fitted  with  a  two-holed  stopper  containing  entrance  and  exit 
tubes,  and  place  this  in  a  beaker  of  cold  water.  Allow  the 
hydrogen  to  pass  through  the  apparatus  for  a  few  moments, 
to  drive  out  the  air ;  then,  when  it  is  safe,  heat  the  copper 
oxid.  Notice  that,  after  action  begins,  the  oxid  glows  brightly. 


ELEMENTS  AND  COMPOUNDS  2/ 

When  all  is  finished,  allow  the  contents  of  the  tube  to  cool 
in  the  stream  of  hydrogen.  Meanwhile  disconnect  the  test 
tube,  and  note  the  water  collected  in  it.  WThen  the  combus- 
tion tube  is  cool,  remove  the  contents  and  recognize  as  me- 


tallic  copper.  Such  a  process  as  this,  i.e.,  the  taking  away 
of  oxygen  from  a  substance,  is  called  reduction.  Adding 
oxygen  to  a  substance  is  called  oxidation.  Which  substance 
in  this  case  is  reduced  ?  Which  oxidized  ? 

Copper  Hydrogen 

4-  Hydrogen  =  „  J  +  Copper 

Oxygen  ^    *  Oxygen 


EXPERIMENT    17 
Zinc  and  Magnesium  Oxids  and  Sulfuric  Acid 

Review  Experiment  15  b,  c,  </,  e,  andy! 

a.  Zinc  Oxid  and  Sulfuric  Acid.  Add  zinc  oxid  to  dilute 
sulfuric  acid  (i  part  acid  to  5  of  water)  in  a  flask,  until  there 
is  no  further  action.  Notice  that  no  hydrogen  is  given  off. 
Evaporate  the  solution,  crystallize,  and  recognize  the  crys- 
tals as  the  same  as  those  obtained  in  Experiment  15  e,  i.e., 
zinc  sulfate.  Since  no  hydrogen  is  given  off,  we  are  com- 


28          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

pelled  to  conclude  that  it  united  with  the  oxygen  of  the  zinc 
oxid,  and  formed  water. 

Zinc          Hydrogen     Zinc 

Oxygen +  «""«       =  Suite  + 
Oxygen         Oxygen 

b.  Magnesium  Oxid  and  Sulfuric  Acid.  Perform  a  similar 
experiment,  using  magnesium  oxid  instead  of  zinc  oxid. 
Name  the  crystals. 

Hydrogen     Magnesium 
Magnesium  ur       =  Sulfur  Hydrogen 

oxygen         Oxygen 


EXPERIMENT    18 
Neutralization 

a.  Sodium  Hydroxid  and  Sulfuric  Acid.  Take  about  5  cc. 
of  concentrated  sulfuric  acid  diluted  with  about  10  cc.  of 
water.  Dissolve  about  10  grms.  of  sodium  hydroxid  in  water. 
Add  the  hydroxid  to  the  acid,  until  the  solution  shows  neither 
an  alkaline  nor  acid  reaction  upon  litmus  paper.  Evaporate 
and  crystallize  it.  Do  the  crystals  contain  water  of  crystalli- 
zation ?  (See  Experiment  15  e.)  Allow  a  dry  crystal  to  lie 
in  the  air  for  some  time,  and  examine  it.  What  has  hap- 
pened to  it  ?  This  phenomenon  is  called  efflorescence,  and 
the  crystal  is  said  to  effloresce. 

Since  sodium  hydroxid  is  composed  of  sodium,  oxygen, 
and  hydrogen,  and  sulfuric  acid  is  composed  of  hydrogen, 
sulfur,  and  oxygen,  and  since  we  have  learned  (see  Ex- 
periment 15  e,  /)  that  the  tendency  of  metals  is  to  force 
hydrogen  out  of  compounds,  it  is  evident  that  the  sodium  of 
the  sodium  hydroxid  has  united  with  the  sulfur  and  oxygen 


ELEMENTS  AND  COMPOUNDS  2Q 

of  the  acid,  leaving  the  hydrogen  of  the  acid  to  unite  with 
the  oxygen  of  the  hydroxid.  The  resulting  crystals  we  call 
sodium  sulfate. 

Sodium        Hydrogen     Sodium      Hydrogen 
Oxygen     +  Sulfur       =  Sulfur    +0xygen 
Hydrogen     Oxygen         Oxygen 

b.  Potassium  Hydroxid  and  Sulfuric  Acid.  Neutralize  po- 
tassium hydroxid  with  sulfuric  acid  as  in  a.  Evaporate  and 
crystallize.  We  name  these  crystals  potassium  sulfate. 

Potassium     Hydrogen     Potassium 
Oxygen      +  Sulfur       =  Sulfur 
Hydrogen      Oxygen         Oxygen 

Remark.  A  compound  obtained  by  replacing  the  hydro- 
gen of  an  acid  by  a  metal,  we  call  a  salt.  We  may  now 
remodel  our  definitions  of  acids  and  bases.  Our  full 
definitions  will  then  be  :  — 

An  acid  is  a  compound  of  one  or  more  non-metallic  ele- 
ments with  hydrogen,  or  with  hydrogen  and  oxygen,  all  or 
part  of  whose  hydrogen  may  be  replaced  by  a  metal. 

A  base  is  a  compound  of  a  metallic  element  with  oxygen, 
or  with  oxygen  and  hydrogen,  the  metal  of  which  is  capable 
of  replacing  the  hydrogen  of  an  acid. 

We  see  that,  in  neutralizing  an  acid  with  a  base,  we  shall 
always  obtain  a  salt  and  water. 

Acidj  Normal,  and  Basic  Salts.  An  acid  salt  is  a  salt 
formed  by  replacing  only  part  of  the  hydrogen  of  an  acid 
by  the  metal  of  a  base. 

A  normal  salt  is  a  salt  formed  by  replacing  all  the  replace- 
able hydrogen  of  an  acid  by  a  metal  of  a  base. 

A  basic  salt  is  a  salt  formed  by  replacing  all  the  re- 
placeable hydrogen  of  an  acid  by  a  metal  of  a  base,  and,  in 
addition,  causing  more  of  the  base  to  unite  with  the  com- 
pound. 


3O        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

EXPERIMENT    19 

Carbonic  Acid  and  Carbon  Dioxid  with  Sodium  and 
Potassium  Hydroxid 

a.  Dissolve  about  i  grm.  of  sodium  hydroxid  in  20  cc.  of 
water,  and  add  from  this  to  a  solution  of  carbonic  acid,  until 
it  no  longer  turns  blue  litmus   paper  red.     Evaporate  and 
crystallize.      Do   the    crystals    contain  water  of    crystalliza- 
tion ?     Do  they  effloresce  ? 

Arguing  as  in  Experiment  18  a,  we  come  to  the  conclu- 
sion that  the  salt  formed  is  made  up  of  the  elements  sodium, 
carbon,  and  oxygen,  chemically  combined.  The  crystals  are 
called  sodium  carbonate  or  sal  soda. 

Sodium  Hydrogen  Sodium 
Oxygen  +  Carbon  =  Carbon 
Hydrogen  Oxygen  Oxygen 

b.  Perform    a   similar    experiment,    using    potassium    hy- 
droxid instead  of  sodium  hydroxid. 

Potassium     Hydrogen      Potassium 

Hydrogen 
Oxygen     +  Carbon     =  Carbon       + 

Hydrogen       Oxygen         Oxygen 

c.  Pass  carbon  dioxid   through  solutions  of  sodium  and 
potassium  hydroxids,  and  obtain  the  same  salts. 

EXPERIMENT    20 
Calcium 

On  account  of  its  expense  and  the  difficulty  of  preserving 
the  element  calcium,  it  is  improbable  that  there  will  be  an 
opportunity  to  notice  its  properties.  Suffice  it  to  say  that  it 


ELEMENTS  AND  COMPOUNDS  31 

is  a  yellowish  metal,  which  may  be  kept  for  some  time  in 
dry  air,  but  which  in  the  presence  of  water,  oxidizes  with  the 
evolution  of  hydrogen,  forming  calcium  hydroxid.  While 
the  element  itself  is  unimportant,  its  compounds  must  be 
carefully  studied.  The  oxid  we  know  familiarly  as  quick 
lime. 

a.  Properties  of  Calcium   Oxid.      Examine  a  lump  of  cal- 
cium oxid,  and  note  its  properties. 

b.  Calcium  Hydroxid.     Place  a  lump  of  calcium  oxid  about 
the  size  of  a  walnut  in   a  porcelain   dish.     Allow  water  to 
drip  upon  this  as  long  as  any  is  absorbed.      Notice  the  evo- 
lution of  heat.     Test  the  product  with  moist  litmus  paper. 

The  water  must  have  united  with  the  calcium  oxid.  thus 
forming  cakium  hydroxid;  just  so  in  Experiment  lor,  water 
united  with  sodium  oxid  and  formed  sodium  hydroxid. 

Calcium 
Calcium          Hydrogen 

=  Oxygen 
Oxygen  Oxygen 

Hydrogen 

c.  Lime  Water.     Put  a  little  of  the  hydroxid  made  in  b 
in  a  beaker  of  cold  water.     Stir,  and,  after  allowing  it  to  set- 
tle, pour  off  the  liquid.     Fill  the  beaker  again  with  water, 
and,  after  stirring  vigorously,  filter  it  into  a  bottle  that  can  be 
tightly  stoppered.    Place  a  little  of  the  liquid  on  a  watch  glass, 
and  evaporate  it.     Is  calcium   hydroxid   soluble  in  water  ? 
The  liquid  in  the  bottle  is  commonly  called  lime  water. 

d.  Lime  Water  and  Carbonic  Acid.      In  a  test  tube,  place 
i   cc.  of  lime  water.     Add  i  cc.  of  carbonic  acid,  and  note 
the  precipitate  formed.     Now  add  more  carbonic  acid,  until 
the    solution  becomes  clear.     Boil   until  the  precipitate  re- 
turns. 

From  a  generator  (described  in  Experiment  26)  pass  car- 
bon dioxid  into  i  cc.  of  lime  water,  until  the  precipitate 
which  at  first  appears  is  dissolved.  Boil  as  before. 


32        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

We  obtain  a  salt  just  as  in  the  experiment  with  sodium 
hydroxid  and  carbonic  acid  ;  but,  in  this  case,  the  salt  is  cal- 
cium carbonate.  This  salt  is  evidently  soluble  in  excess  of 
carbonic  acid,  and  reappears  in  the  solution  when  the  car- 
bonic acid  is  decomposed  by  heat  (see  Experiment  8  <•). 
This  action,  then,  can  be  used  as  a  test  for  the  presence  of 
carbon  dioxid,  or,  conversely,  as  a  test  for  the  presence  of 
calcium  hydroxid.  See  whether  or  not  your  breath  *  con- 
tains carbon  dioxid. 

Calcium        Hydrogen     Calcium 
(0          Oxygen     +  Carbon      =  Carbon   + 
Hydrogen     Oxygen         Oxygen 

Calcium  Calcium 

i^\  Carbon  Hydrogen 

Oxygen     +rt  =  Carbon   +    * 

Oxygen  Oxygen 

Hydrogen  Oxygen 


EXPERIMENT    21 
Analysis  of  Marble  (Calcium  Carbonate) 

Take  a  piece  of  marble  about  the  size 
of  a  hazel  nut,  and   grind   it  to   a   fine 
^>    ^Lr  powder  in  a  mortar.     Place  this   in  a 

hard  glass  tube  closed  at  one  end.  Fit 
the  tube  with  a  one-holed  stopper  con- 
taining an  exit  tube  extending  a  centi- 
meter beneath  the  surface  of  a  little 
lime  water  in  a  test  tube.  Heat  the 

*  The  production  of  carbon  dioxid  by  the  processes  of  respiration 
and  combustion  depletes  the  supply  of  oxygen  in  a  room.  The  oxygen 
must  be  renewed  from  the  air  outside,  and  the  injurious  respiration 
and  combustion  products  must  be  removed  by  proper  ventilation.  This 
is  especially  true  of  churches,  theaters,  schoolrooms,  and  all  in-door 
places  where  many  persons  are  assembled  at  the  same  time. 


ELEMENTS    AND    COMPOUNDS 


33 


hard  glass  tube  with  a  blast  lamp.  After  the  air  has 
been  driven  out,  notice  the  effect  of  the  bubbles  upon  the 
lime  water.  Continue  heating  it  until  the  milkiness  of  the 
lime  water  disappears.  Then  take  away  the  lime-water  tube 
(before  taking  away  the  lamp),  and  heat  it  over  a  Bunsen 
burner.  Does  the  milkiness  return  ?  We  therefore  know 
that  one  constituent  of  marble  is  carbon  dioxid  (see  Experi- 
ment 20  d). 

Take  another  lump  of 
marble  about  the  size  of  a 
pea,  and,  holding  with  a  pair 
of  forceps  over  a  porcelain 
plate,  heat  it  with  a  blast 
lamp  for  about  five  minutes. 
Collect  the  particles  that 
have  fallen  on  the  plate, 
and.  together  with  what  is 
left  of  the  lump,  place  them 
in  a  very  small  quantity  of 
water  (only  enough  to  mois- 
ten them)  on  a  watch  glass.  Test  this  with  moist  litmus  and 
turmeric  paper.  Now  add  more  water  (about  i  cc.),  and 
filter  into  a  test  tube.  Add  a  little  carbonic  acid,  and  notice 
the  milkiness  formed.  On  adding  more  carbonic  acid,  note 
that  it  disappears,  and  reappears  on  heating.  Hence  we  see 
that  another  constituent  of  marble  is  calcium  oxid  (see  Ex- 
periment 20  d),  and  that  marble  is  composed  of  calcium, 
carbon,  and  oxvgen.  Its  chemical  name  is  calcium  carbonate. 


Calcium 
Carbon   = 


Calcium 
Oxygen 


Carbon 
Oxygen 


Oxygen 

To  make   a  complete  analysis,  we    must  take  two  more 
steps.     First  take  a  jar  of  carbon  dioxid,  and,  after  igniting 


34        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

a  ribbon  of  magnesium,  plunge  it  into  the  jar.  Note  that 
the  magnesium  burns  to  the  white  oxid,  and  that  carbon  is 
deposited  on  the  sides  of  the  jar. 

Carbon    ,  Magnesium 

+  Magnesium  ==  4-  Carbon 

Oxygen T  Oxygen 

Secondly  take  about  5  grins,  of  calcium  oxid,  and  grind  it 
to  a  powder  in  a  mortar.  Mix  with  this  2.5  grms.  of  metallic 
magnesium  powder,  place  the  mixture  in  an  iron  crucible, 
and  heat  it  over  a  Bunsen  burner.  While  the  contents  of  the 
crucible  are  still  warm,  hold  it  by  means  of  a  pair  of  long 
forceps  under  a  beaker  full  of  water  and  inverted  on  the 
shelf  of  a  pneumatic  trough.  Recognize  the  gas  collected 
as  hydrogen. 

Since  neither  magnesium,  magnesium  oxid,  nor  calcium 
oxid  decomposes  water  with  the  evolution  of  hydrogen  (see 
Experiments  13  and  20),  it  must  be  that  metallic  calcium  was 
formed  in  the  crucible  and,  on  being  placed  in  the  water, 
decomposed  it,  giving  hydrogen. 

/  N        Calcium  Magnesium 

1 i )  _  +  Magnesium  =  _    '  -f  Calcium 
Oxygen                             Oxygen 

Calcium 

(2)  Calcium  +    y  r°gen     =  Oxygen        +  Hydrogen 

xygen  Hydrogen 


EXPERIMENT    22 
Chlorin 

a.  Properties.  Apply  to  the  instructor  for  a  jar  of  chlorin. 
In  working  with  this  element,  perform  your  experiments  under 
a  hood  with  a  strong  draught,  since  chlorin  taken  into  the 
lungs  is  very  dangerous.  Note  the  color  and  odor  (you  can 


ELEMENTS    AND    COMPOUNDS 


35 


hardly  avoid  smelling  it).  Moisten  a  little  blue  and  red 
litmus  paper,  also  a  colored  flower,  if  possible,  and  allow 
them  to  hang  in  the  gas  for  a  few  moments.  Pour  about 
100  cc.  of  water  into  the  jar  ;  seal,  shake,  and  open  it  under 
water.  Is  chlorin  soluble  in  water  ? 

b.  Hydrogen  Chlorid  (Hydrochloric  Acid)  by  Burning  Hy- 
drogen in  Chlorin.  Have  ready  a  hydrogen  generator  (see 
Experiment  16  a\  When  the  air  is  driven  out,  ignite  the 
jet  and  allow  it  to  burn  in  a  jar  of  chlorin  until  no  green 


color  remains.  Blow  your  breath  across  the  mouth  of  the 
jar,  and  notice  that  the  gas  fumes  in  the  presence  of  mois- 
ture. Test  with  moist  litmus  paper.  Taste  the  gas  by 
allowing  a  little  to  enter  the  mouth.  This  is  called  hydrogen 
chlorid,  or  hydrochloric  acid  gas. 


c.  Union  of  Hydrogen  and  Chlorin  by  means  of  Light.  To 
be  performed  by  the  instructor.  In  a  dark  room  with  ruby 
light  (a  photographic  dark-room),  fill  a  250  cc.  flask  one- 
half  full  of  chlorin  and  one-half  full  of  hydrogen,  collected 
over  warm  water.  Fit  a  rubber  stopper  to  the  flask,  and 


36        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

fasten  it  securely  with  wire.  Wrap  a  towel  or  thick  cloth 
around  the  flask  so  that  no  light  can  enter.  Tie  a  long 
string  to  a  corner  of  the  towel,  and  carry  all  out  of  doors.  At 
a  safe  distance,  drag  off  the  towel  by  means  of  the  string  so 
that  the  sunlight  may  strike  the  flask.  What  caused  the 
action,  and  what  is  formed  ? 

d.  Sodium  Chlorid.  Apply  to  the  instructor  for  a  jar  full 
of  dry  chlorin.  Into  this,  drop  a  thin  slice  of  metallic  sodium 
about  half  a  square  centimeter  in  area,  and  seal.  Allow  it 
to  stand  until  the  next  day.  Then  open  the  jar,  and  remove 
the  white  compound.  Crush  it,  be  sure  there  is  no  metallic 
sodium  left,  and  then  taste  it.  What  is  it  ?  Being  formed 
by  the  synthesis  of  sodium  and  chlorin,  it  is  called  sodium 
chlorid.  Note  that,  in  cases  where  two  elements  unite,  the 
name  always  ends  in  id. 

Sodium 
Sodium  +  Chlorin  =  _.  , 

Chlorin 


EXPERIMENT   23 
Hydrochloric  Acid  from  Sodium  Chlorid  and  Sulf uric  Acid 

In  a  250  cc.  flask  fitted  with  a  two-holed  stopper  contain- 
ing funnel  and  exit  tubes,  place  about  10  grins,  of  sodium 
chlorid.  Have  the  exit  tubes  connected  with  a  glass  tube 
reaching  to  the  bottom  of  a  fruit  jar  covered  with  a  piece  of 
cardboard.  Into  the  funnel  tube,  pour  15  cc.  of  concen- 
trated sulfuric  acid  diluted  with  4  cc.  of  water.  Pour  the 
acid  into  the  water.  Collect  two  or  three  jars  of  the  gas. 
Note  that  it  is  the  same  gas  as  that  obtained  by  burning  hy- 
drogen in  chlorin.  To  one  of  the  jars,  add  about  25  cc.  of 
water,  close  and  shake  it.  Open  it  under  water.  Is  hydro- 
chloric acid  soluble  in  water  ?  Into  another  jar,  pour  a  little 


ELEMENTS  AND  COMPOUNDS  37 

water,  shake  and  test  ttie  liquid  with  blue  litmus  paper.  Di- 
lute the  contents  of  the  flask  with  water  after  all  the  gas  pos- 
sible has  been  allowed  to  escape.  Evaporate  and  crystallize. 
Recognize  the  crystals  as  the  same  as  those  obtained  in  Ex- 
periment 1  8  a,  i.e.,  sodium  sulfate. 

Hydrogen      Sodium 
Sodium       _'  Hydrogen 


Chlorin  Chlorin 

Oxygen          Oxygen 


EXPERIMENT    24 
Analysis  of  Hydrochloric  Acid 

a.  By  Electrolysis.     Have  ready  an  apparatus  like  that  of 
Experiment  6  e,  except  that  the  electrodes  are  of  gas  carbon 
instead  of  platinum.      Fill  it  with  concentrated  hydrochloric 
acid  solution.    On  turning  on  the  current,  note  the  immediate 
evolution  of  hydrogen  at  the  negative  pole.     Since  you  know 
that  chlorin  is  soluble  in  water,  you  can  easily  account  for 
the  fact  that  chlorin  does  not   immediately  appear    at   the 
positive  pole.     Allow  the  current  to  pass  through  the  solu- 
tion for  some  time  ;  then,  when  it  has  taken  up  all  the  chlorin 
it  can  hold,  hydrogen   and  chlorin  will  appear  in  equal  vol- 
umes at  the  negative  and  positive  poles  respectively. 

Hydrogen  CMor.n 

Chlorin 

b.  By  Means  of  Sodium.     Fill  a  dry  quick-sealing  fruit  jar 
with  hydrochloric  acid  gas  dried  by  being  passed  through 
concentrated  sulfuric  acid.     Be    sure    to    have    the    rubber 
washer  well  greased  with  vaseline,  so  that  the  jar  will  be  air 
tight  when  closed.     Prepare  about  25  grms.  of  sodium  amal- 
gam (10  parts  mercury  to  i  part  sodium),  and  powder  it  in  a 


38         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

mortar.  Drop  the  powdered  amalgam  into  the  jar,  seal 
quickly,  shake  vigorously,  and  note  the  action.  Open  it  under 
water  in  a  glass  or  porcelain  dish,  and,  allowing  the  mercury 
to  fall  out,  note  the  volume  of  the  remaining  gas.  Remove 
the  jar  from  the  water,  and  recognize  the  gas  as  hydrogen. 
If  the  experiment  is  successful,  you  will  have  one-half  a  jar 
of  hydrogen,  showing  that  hydrochloric  acid  is  composed  of 
equal  volumes  of  hydrogen  and  chlorin. 


EXPERIMENT    25 

Hydrochloric  Acid  with  Metals,  Hydroxids,  and 
Carbonates 

a.  Try  the  effect  of  hydrochloric  acid  on  the  following 
metal.5  that  you  have  studied,  —  iron,  magnesium,  zinc,  and 
copper,  and  verify  the  following  statements  :  — 

Hydrogen     Iron 

(1)  Iron+_;\  -  +  Hydrogen 

Chlorin         Chlorin 

Hydrogen      Magnesium 

(2)  Magnesium  +  =  +  Hydrogen 


Hydrogen      Zinc 
Chlorin     =  Chlorin 

Name  the  salts  formed. 

b.  Neutralize  a  small  amount  of  hydrochloric  acid,  first 
with  sodium  hydroxid,  second  with  potassium  hydroxid, 
and  third  with  sodium  carbonate,  and  verify  the  following 
statements  :  — 

,  ^  Hydrogen  _  Sodium     Hydrogen 

W        Oxygen     +  Chlorin     ~  chlorin  +  Oxygen 
Hydrogen 


ELEMENTS    AND    COMPOUNDS  39 

Potassium 

Hydrogen     Potassium      Hydrogen 
(2)        Oxygen      +Chlorin     =  chlorin      + 

Hydrogen 

.  Hydrogen  _  Sodium     Carbon      Hydrogen 

+  Chlorin     ~  Chlorin  +  Oxygen     Oxygen 
Oxygen 

Name  the  salts  formed. 

c.  Potassium  Chlorid.  Compare  the  salt  formed  in  b  i 
by  neutralizing  potassium  hydroxid  and  hydrochloric  acid, 
with  the  residue  obtained  in  the  hard  glass  tube  in  Experi- 
ment 6  c.  We  now  see  that  potassium  chlorate  is  composed 
of  potassium,  chlorin,  and  oxygen,  and  that  the  action  in 
Experiment  6  c,  can  be  expressed  by  the  following :  — 

Potassium     Potassium 

Chlonn       =  Chlorin       +0xysen 
Oxygen 

EXPERIMENT    26 
Carbon  Dioxid  from  Marble  and  Hydrochloric  Acid 

The  best  way  to  obtain  a  stream  of  carbon  dioxid  for  use 
in  the  laboratory  is  by  the  action  of  hydrochloric  acid  upon 
marble  (calcium  carbonate).  In  a  500  cc.  flask  fitted- with 
funnel  and  exit  tubes,  place  a  number  of  lumps  of  marble, 
and  then  add  dilute  hydrochloric  acid  (i  part  acid  to  i  part 
water).  Carbonic  acid  is  first  formed,  which,  being  very  un- 
stable (see  Experiment  8  c],  breaks  up  into  carbon  dioxid 
and  water. 

<;al<iium     Hydrogen     Calcium      ^f"^ 

(1)  Carbon    +^J      =  chlorin   -Carbon 

Oxygen  Oxygen 

Hydrogen 

(2)  Carbon      =  Carbon      Hydrogen 

Oxygen        Oxygen     Oxygen 


4O         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


EXPERIMENT    27 
Preparation  of  Chlorin 

a.  Test  for  Manganese.  Take  a  piece  of  platinum  wire,  and 
make  a  loop  on  one  end  about  3  mm.  in  diameter.  Heat  it 
in  the  Bunsen  flame,  and,  while  it  is  still  hot,  dip  it  into  a 
little  microcosmic  salt.  Place  it  in  the  flame  again,  and  make 
a  bead  of  the  fused  salt.  Melt  into  it  a  grain  of  powdered 
manganese  dioxid,  and,  after  holding  in  the  flame  for  a  few 


moments,  remove  it,  and   note  the   characteristic  amethyst 

color.     Try  other  manganese  compounds  in  the  same  manner. 

b.    Preparation.     Under  a  hood,  place  a  flask  containing 

six  or   seven  small   lumps  of  manganese  dioxid,   and   add 


ELEMENTS    AND    COMPOUNDS  4! 

enough  hydrochloric  acid  to  cover  them.  Warm  the  flask, 
and  recognize  the  gas  evolved  as  chlorin.  (See  illustration 
on  page  40.)  This  is  the  way  your  instructor  prepared  the 
gas  for  you.  After  as  much  gas  as  possible  has  been 
evolved,  filter  the  liquid,  evaporate,  and  crystallize.  Put  a 
crystal  or  two  in  a  test  tube,  and  add  a  little  concentrated 
sulfuric  acid.  Notice  that  hydrochloric  acid  is  evolved. 
Therefore,  since  hydrochloric  acid  contains  chlorin,  and, 
since  that  chlorin  could  not  have  come  from  the  sulfuric 
acid,  it  must  have  come  from  the  crystals.  Make  a  micro- 
cosmic  bead  on  a  loop  of  platinum  wire  as  in  a,  add  a  little 
of  the  crystals,  and  heat  in  the  hottest  part  of  the  flame. 
The  characteristic  amethyst  color  proves  the  presence  of 
manganese  in  the  crystals.  It  is  evident,  therefore,  that  the 
crystals  are  manganese  chlorid.  The  only  way  we  can 
account  for  the  oxygen  of  the  manganese  dioxid  is  that  it 
united  with  the  hydrogen  of  the  hydrochloric  acid  and 
formed  water.  Therefore  we  have  the  following :  — 

Manganese      Hydrogen  _  Manganese      Hydrogen  . 

Oxygen  Chlorin     ~~  Chlorin  Oxygen 

EXPERIMENT    28 
Bromin 

a.  Properties.  Examine  under  a  hood  a  small  amount  of 
bromin,  and  note  its  properties,  —  color,  odor,  weight,  etc. 
Be  careful  not  to  inhale  the  fumes,  since  bromin  acts  vio- 
lently upon  the  membranes  of  the  throat  and  lungs.  A  drop 
of  bromin  on  the  skin  produces  a  severe  wound.  Note  its 
solubility  in  water,  alcohol,  ether,  and  carbon  bisulfid.  Have 
no  fire  near,  since  alcohol,  ether,  and  carbon  bisulfid  are 
very  inflammable.  Hang  moist  pieces  of  litmus  paper  in 
vapors  of  bromin,  and  note  the  bleaching  action. 


42         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

b.  Hydrogen  Bromid  (Hydrobromic  Acid).  Have  ready 
under  a  .hood  a  test  tube  containing  about  i  cc.  of  bromin. 
Fit  this  with  a  two-holed  stopper  containing  an  entrance  and 
an  exit  tube.  Let  the  entrance  tube  extend  below  the  sur- 
face of  the  bromin.  Attach  the  entrance  tube  to  a  hydro- 
gen generator,  so  that  the  hydrogen  must  bubble  through 
concentrated  sulfuric  acid  in  a  catch  bottle. 


Attach  the  exit  tube  to  a  hard  glass  tube  about  20  cm. 
long,  containing  platinized  asbestos.  Fit  a  rubber  tube 
about  20  cm.  long  to  this.  Allow  the  hydrogen  to  pass 
through  the  apparatus  until  all  the  air  is  expelled.  Then 
heat  the  platinized  asbestos  to  redness,  warming  the  bromin 
very  gently.  Collect  the  new  gas  in  a  jar  by  displacement 
of  air,  and  get  its  properties.  This  gas  is  called  hydrogen 
broinid,  or  hydrobromic  acid. 

Hydrogen  +  Bromin  =  Hydrogen 
Bromin 

c.  Hydrobromic  Acid  and  Potassium  Hydroxid,  Neutral- 
ize a  solution  of  hydrobromic  acid  with  potassium  hydroxid, 
and  verify  the  following  :  — 

Potassium 

Hydrogen  _  Potassium     Hydrogen 

_   ,  Bromin     ~  Bromin          Oxygen 

Hydrogen  7* 


ELEMENTS  AND  COMPOUNDS  43 

d.  Replacement  of  Bromin   by    Chlorin.      Dissolve   about 
a  gram  of  sodium   or    potassium  bromid  in   water.      After 
adding  a  little  chlorin  water   and  about    i    cc.  of  ether  or 
carbon  bisulfid,  shake  it  vigorously.     Which  is  the  stronger 
chemically,  chlorin  or  bromin  ? 

Potassium  Potassium 

+  Chlorin  =  _,  ,  +  Bromin 

Bromin  Chlorin 

e.  Potassium    Bromid  and  Sulfuric  Acid.      Try  to  make 
hydrobromic     acid,     using    sulfuric     acid    and     potassium 
bromid. 


EXPERIMENT   29 
lodin 

a.  Properties.    Take  a  few  crystals  of  iodin,  examine,  and 
note  properties.     Warm  a  small  crystal  in  a  test  tube,  and 
note   the    color,    odor,    and    specific   gravity  of   the  vapor. 
Test  the  solubility   of  iodin   in  water,    alcohol,   ether,  and 
carbon  bisulfid.     Take  a  small  lump  of  starch,  add  a  few 
drops  of  water  and  mix  into  a  paste.     Then  heat  a  test  tube 
of  water  to  boiling,  and   add  this    quickly  to   the   starch. 
Place  a  little  of  the  starch  solution  in  a  test  tube  of  water, 
shake  it,  and  add  a  little  of  the  water  solution  of  iodin.    The 
blue  color  obtained  is  a  characteristic  test  for  the  presence 
of  uncombined  iodiri. 

b.  Phosphorus  lodid.     Weigh  out  accurately  about  i  grm. 
of  phosphorus,    and  dissolve   it  in    about  5   cc.   of   carbon 
bisulfid  in  a  test  tube.     Then  weigh  out  8.2  times  as  much 
iodin.      Add   the    iodin  gradually  to  the  phosphorus   solu- 
tion.    When  the  iodin  has  disappeared,  drive  off  the  carbon 
bfsulfid  by  placing  the  test  tube  in  warm  water.     Do  this 


44         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

under  a  hood.     The  resulting  crystals  must  be  phosphorus 
iodid.* 

Phosphorus 

Phosphorus  +  lodm  =  _  ,. 
loam 

Hydriodic  Acid.  By  passing  vapors  of  iodin  and  hydro- 
gen over  finely  divided  platinum  raised  to  a  red  heat,  the 
two  elements  may  be  made  to  combine,  but  not  readily. 
When  made  by  this  or  other  methods,  hydrogen  iodid  is 
a  colorless,  heavy  gas,  very  soluble  in  water,  fumes  in  moist 
air,  and  gives  an  acid  reaction  with  litmus  paper.  It  is  very 
similar  to  hydrochloric  and  hydrobromic  acids. 


EXPERIMENT  30 
Potassium  Iodid 

a.  Hydriodic  Acid  and  Potassium  Hydroxid.  Neutralize  a 
small  quantity  of  dilute  hydriodic  acid  with  potassium  hy- 
droxid,  and  verify  the  following  :  — 

Hydrogen     P°tassium     Potassium  ,  Hydrogen 

+       =Ioain 


b.  Potassium  Iodid  and  Sulf  uric  Acid.    In  a  test  tube,  place 
a  crystal    of   potassium   iodid.      Add    a   little   concentrated 
sulfuric  acid.     Do  you  get  hydriodic  acid  ? 

c.  Potassium  Iodid  and  Phosphoric  Acid.     In  a  test  tube, 
place  a  crystal  of  potassium  iodid,  and  add  a  little   strong 
phosphoric  acid.     Heat  it,  and  see  whether  you  obtain  hydri- 
odic acid. 

*  When  phosphorus  iodid  is  thrown  into  water,  hydriodic  acid  is 
formed,  part  of  the  hydrogen  of  the  water  uniting  with  the  iodin  of 
the  iodid,  * 


ELEMENTS  AND  COMPOUNDS  45 

d.  Potassium  lodid  and  Starch  Paste.     Try  the   effect  of 
starch  paste  upon  potassium  iodid. 

e.  Solubility  of  lodin  in  Potassium  Iodid  Solution.     Dissolve 
a  little  potassium  iodid  in  water  in  a  test  tube.     Add  a  crys- 
tal of  iodin.     Is  iodin  soluble  in  potassium  iodid  ? 

f.  Replacement  of  Iodin  by  Chlorin  and  Bromin.  Dissolve 
a  little  potassium  iodid  in  water  in  a  test  tube.  Add  chlorin 
water,  and  separate  into  two  equal  parts.  To  one  add  starch 
paste,  to  the  other  add  carbon  bisulfid,  and  shake  them. 
Which  is  the  stronger  chemically,  chlorin  or  iodin  ? 

Potassium  ,  Potassium 

_  +  Chlorin  =  „. .  +  Iodin 

Iodin  Chlorin 

Try  the  same,  using  bromin  water  instead  of  chlorin  water. 
Which  is  the  stronger  chemically,  bromin  or  iodin  ? 

Potassium  Potassium 

_  ,.  +  Bromin  =  _        .         +  Iodin 

Iodin  Bromin 

Fluorin. 

There  is  an  element  fluorin,  a  colorless  gas,  which  in  a 
great  many  ways  is  similar  to  chlorin,  bromin,  and  iodin. 
It  is  the  most  active  chemically  of  all  the  elements,  and 
was  not  known  in  the  uncombined  state  until  comparatively 
recently.  One  great  difficulty  in  preparing  it  was  to  find 
material  for  vessels  that  it  would  not  attack.  Its  principal 
salt  is  calcium  fluorid,  which  occurs  in  nature  as  the  mineral 
fluor  spar. 

EXPERIMENT    31 
Calcium  Fluorid  and  Hydrogen  Fluorid 

a.  Properties  of  Calcium  Fluorid.     Examine  calcium  fluorid 
both  in  crystal  and  powder  form. 

b.  Hydrogen  Fluorid  (Hydrofluoric  Acid).     In  a  test  tube, 


46 


AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


put  i  cc.  of  powdered  calcium  fluorid.  Cover  with  sul- 
furic  acid,  and  heat  it  under  the  hood.  Get  the  color  and 
odor  of  the  gas,  being  careful  to  smell  cautiously,  since  the 
gas  seriously  affects  the  membranes  of  the  throat  and  lungs. 
Breathe  over  the  mouth  of  the  test  tube.  Test  the  gas  with 
blue  litmus  paper.  Remove  the  contents  of  the  tube,  wash 
and  dry  it.  Examine  the  glass,  noticing  how  it  has  been 
corroded.  This  gas  is  called  hydrofluoric  acid.  Assuming 
that  calcium  fluorid  is  composed  of  calcium  and  fluorin,  we 
readily  see  that 


Calcium 
Fluorin 


Hydrogen 
+  Sulfur       = 


_  Hydrogen 
Fluorin 


Calcium 

Sulfur 

Oxygen 


Oxygen 

c.    Etching  Glass.     In  a  shallow  lead  dish,  make  a  paste  cf 
powdered  calcium   fluorid    and  concentrated    sulfuric   acid. 

f_ _     Take  about   half    as  many 

cc.  of  acid  as  you  take 
grins,  of  calcium  fluorid. 
Cover  a  piece  of  window- 
glass  with  a  thin  layer  of 
melted  paraffin,  warming 
the  glass  to  spread  the  par- 
affin evenly.  When  the 
glass  is  cool,  trace  a  figure 
on  it,  and,  laying  it  upon 
the  dish,  leave  it  overnight. 
Glass  is  composed  of 
various  substances,  among 
which  is  the  element  silicon. 
Fluorin  has  a  strong  affinity  for  this  element ;  hence  it  de- 
composes the  glass  by  abstracting  the  silicon,  forming  a 
colorless  gas,  silicon  fluorid.  Thus  the  hydrofluoric  acid  cor- 
rodes the  glass  where  it  was  not  protected  by  the  paraffin. 


ELEMENTS  AND  COMPOUNDS 


47 


EXPERIMENT    32 
Sulfids 

a.  Synthesis  of  Hydrogen  Sulfid.  Have  ready  a  hydrogen 
generator.  Take  a  test  tube,  place  in  it  about  i  o  grms.  of 
sulfur,  and  fit  it  with  a  two-holed  stopper  containing  an 
entrance  tube  leading  within  2  cm.  of  the  sulfur,  and  an  exit 
tube  bent  at  right  angles  twice,  the  second  bend  extending 


upwards  about  i  o  cm.  and  drawn  to  an  opening  of  about 
i  mm.  diameter.  Connect  with  the  hydrogen  generator 
so  that  the  hydrogen  must  pass  through  a  catch  bottle 
containing  concentrated  sulfuric  acid,  to  dry  it.  Allow  the 
hydrogen  to  pass  through  the  apparatus  in  order  to  drive 
out  the  air  ;  then,  when  all  is  safe,  heat  the  sulfur  to  boiling. 
Note  the  odor  of  the  escaping  gas.  Light  it,  and  get  the 
color  of  the  flame.  Heat  the  tube  gently  above  the  stopper, 
and  note  that  sulfur  is  deposited.  Is  hydrogen  strongly 
united  with  sulfur  in  hydrogen  sulfid  ? 


48         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


b.  Sulfids  of  Metals,  i  .  Mix  equal  volumes  of  fine  iron 
filings  and  flowers  of  sulfur.  Heat  the  mixture  in  a  test  tube. 
After  action  has  ceased,  break  open  the  tube  and  get  the 
properties  of  the  resulting  iron  sulfid. 


Iron  +  Sulfur  = 

Sulfur 

2.  Mix  6.5  parts  of  zinc  dust  and  3.2  parts  of  the  flowers 
of  sulfur,  and  place  the  mixture  on  an  iron  plate.  In  mixing, 
do  not  use  pressure,  since  the  two  elements  may  unite  with 
explosive  violence.  Cautiously  light  the  mixture,  and  get  the 
properties  of  the  resulting  zinc  sulfid. 


Zinc  +  Sulfur  = 

Sulfur 

3.  Place  in  a  test  tube  a  mixture  of  6.4  parts  of  copper 
filings  with  3.2  parts  of  flowers  of  sulfur,  and  heat  as  in  i. 
Get  the  properties  of  the  resulting  copper  sulfid. 

In  a  test  tube,  boil  a  small  quantity  of  sulfur.  By  means 
of  a  pair  of  long  forceps,  hold  a  strip  of  copper  in  a  flame 

until  red  hot;  then  quickly 
hold  it  in  the  vapors  of  sulfur 
in  the  tube. 

Copper  +  Sulfur  = 

How  do  these  sulfids  differ 
chemically  from  the  corre- 
sponding sutfates? 

c.  Hydrogen  Sulfid  from 
Iron  Sulfid  and  Sulfuric  Add. 
In  a  100  cc.  flask,  place  about  10  grms.  of  iron  sulfid  in  small 
lumps.  Add  dilute  sulfuric  acid  (3  parts  acid  to  i  part 


ELEMENTS  AND  COMPOUNDS  49 

water),  and  warm  it.  Collect  in  jars  and  recognize  the  same 
gas  as  was  obtained  in  a,  i.e.,  hydrogen  sulfid.  Is  it  soluble 
in  water  ?  Note  its  action  on  blue  litmus  paper.  Ignite  a 
jar  of  the  gas,  and  note  what  the  products  of  the  combustion 
are.  When  all  the  gas  possible  has  been  evolved,  filter, 
evaporate,  and  crystallize  it.  Recognize  the  crystals  as  iron 
sulfate.  (See  Experiment  1 5  /  ) 

Hydrogen  Iron 

Iron          _J '       *          Hydrogen      _   _, 
_   ,,     +  Sulfur       =e  .-  +  Sulfur 

Sulfur  Sulfur 

Oxygen  Oxygen 

If  you  had  used  the  sulfids  made  in  2  and  3.  what  salts 
would  you  have  obtained  instead  of  iron  sulfate  ? 

d.  Action  of  Bromin  upon  Hydrogen  Sit/fid.  In  a  100  cc. 
flask,  place  about  50  cc.  of  water  ;  add  a  drop  of  bromin, 
and  shake  it  until  it  is  dissolved.  Fit  the  flask  with  an  exit 
and  entrance  tube,  and,  under  a  hood,  pass  hydrogen  sulfid 
through  the  liquid.  Note  the  disappearance  of  the  red  color. 
Filter  the  solution ;  then  recognize  the  residue  as  sulfur  by 
holding  it  with  a  pair  of  forceps  in  a  Bu.nsen  flame. 

Hydrogen     Bromin  =  Hydrogen 
Sulfur  Bromin 


EXPERIMENT    33 
Other  Compounds  of  Carbon 

a.  Carbon  Monoxid.  Fill  a  rubber  gas-bag  with  dry  carbon 
dioxid  made  from  marble  and  hydrochloric  acid  (see  Experi- 
ment 26).  Place  in  a  hard  glass  combustion  tube  about 
5  grms.  of  zinc  dust,  and  clamp  this  to  a  support.  To  one 
end,  attach  an  empty  bag  similar  to  the  first ;  to  the  other 
end,  the  bag  filled  with  the  gas.  Heat  the  tube  with  a  blast 
lamp,  and  pass  the  carbon  dioxid  back  and  forth  over  the  hot 


50         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

zinc.  Notice  the  color  of  the  compound  forming  in  the  tube, 
both  when  it  is  hot  and  cold.  It  evidently  corresponds  to  the 
zinc  oxid  formed  in  Experiment  12  b.  The  only  substance 
from  which  the  zinc  could  have  obtained  so  much  oxygen 
must  be  the  carbon  dioxid.  Has  it  removed  all  the  oxygen 
from  the  gas  ?  If  so,  there  would  be  nothing  but  carbon  left. 
Force  all  the  gas  into  one  bag  and  confine  it  by  means  of  a 


pinch  cock.  In  a  glass  vessel  filled  with  a  moderately  strong 
solution  of  sodium  hydroxid,  invert  a  beaker  filled  with  the 
liquid,  and  collect  some  of  the  gas  from  the  bag.  The  solu- 
tion of  sodium  hydroxid  will  absorb  any  carbon  dioxid  re- 
maining unchanged  (see  Experiment  19  a),  and  will  allow  the 
new  product  to  pass  through.  Get  the  properties  of  the  new 
gas.  Lift  the  beaker  from  the  solution  ;  immediately  ignite 
the  gas  with  a  match,  and  note  the  color  of  the  flame.  This 
new  gas  must  be  composed  of  carbon  and  oxygen,  but  it 
must  have  less  oxygen  than  carbon  dioxid,  since  the  zinc 
removed  a  part.  It  is  called  carbon  monoxid.  Which  sub- 
stance is  reduced  ?  Which  is  oxidized  ? 

Carbon      ^^  _  Zinc  Carbon 

Oxygen  ~~  Oxygen      Oxygen 

b.  Analysis  of  Oxalic  Acid.  In  a  hard  glass  tube  closed  at 
one  end,  and  fitted  with  a  one-holed  stopper  containing  an 
exit  tube,  place  about  5  grms.  of  oxalic  acid  crystals.  Con- 


ELEMENTS    AND    COMPOUNDS 


nect  the  exit  tube  with  a  dry  empty  test  tube  placed  in  cold 
water,  and  with  this  connect  a  third  test  tube  containing  lime 
water,  so  that  any  gas  coming  from  the  hard  glass  tube  must 
pass  to  the  bottom  of  each  test  tube.  Now  heat  the  hard 
glass  tube  very  gently,  having  it  inclined  at  an  angle  of  30°, 
and  collect  the  evolved  gas  over  a  sodium  hydrate  solution,  as 
Recognize  this  gas  as  carbon  monoxid.  Note  the  for- 


m  a. 


mation  of  water  in  the  first  test  tube.     Note  also  the  milki- 
ness  which  is  at  first  produced  in  the  lime  water,  but  which 


soon  disappears,  proving  the  presence  of  carbon  dioxid. 
The  sodium  hydrate  removes  all  traces  of  this  from  the  car- 
bon monoxid.  Therefore,  we  see,  oxalic  acid  is  composed  of 
water,  carbon  dioxid,  and  carbon  monoxid,  i.e.,  carbon,  oxy- 
gen, and  hydrogen. 


Carbon 
Oxygen     = 
Hydrogen 


+  Carbon 
Oxygen 


(monoxid) 

Carbon 

Oxygen 


52         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


c.  Carbon  Monoxid  from    Oxalic  Add.     In  a  250  cc.  flask 
fitted  with  a  funnel  and  exit  tube,  place  about  20  grins,  of 
oxalic  acid  crystals.     Add  through  the  funnel  tube  enough 
concentrated  sulfuric  acid  to  cover  the  crystals.    Heat  it,  and 
pass  the  evolved  gas  first  through  an  empty  catch  bottle, 
and  then  through  a  second  containing  a  strong  solution  of 
sodium  hydrate.     Collect  the  gas  over  water,  and  recognize 
as  carbon  monoxid. 

We  know  that  sulfuric  acid  has  a  strong  attraction  for 
water.  If,  then,  it  takes  water  away  from  oxalic  acid,  we 
shall  have  only  the  two  gases  carbon  dioxid  and  carbon 
monoxid  left.  The  carbon  dioxid  is  absorbed  by  the  sodium 
hydroxid  in  the  catch  bottle. 

d.  Methyl  Hyd rid  (Marsh  Gas)  from  Carbon  Monoxid  and 
Hydrogen.     In  a  eudiometer  tube  fitted  with  two   platinum 

wires  fused  into  the 
closed  end,  collect  over 
mercury  10  cc.  of  carbon 
monoxid  and  30  cc.  of  hy- 
drogen, both  gases  being 
dry.  Allow  electric  sparks 
from  an  induction  coil  to 
pass  between  the  points. 
Notice  the  formation  of 
moisture  and  the  decrease 
in  volume.  Get  the  prop- 
erties of  the  new  gas.  It 
is  called  methyl  hydrid, 
or  marsh  gas,  and  is  com- 
posed of  carbon  and  hy- 
drogen. Part  of  the  hy- 
drogen must  have  united  with  oxygen  to  form  water,  and 
part  must  have  united  with  the  carbon  to  make  the  new 


ELEMENTS  AND  COMPOUNDS  53 

gas  ;    otherwise   only  water    and  carbon    would  have   been 
formed. 

arb0n  ydrogen 


+  Hydrogen  = 

Oxygen  Hydrogen     Oxygen 

e.  Marsh  Gas  from  Sodium  Acetate  and  Soda  Lime.  In  a 
hard  glass  tube,  place  an  intimate  mixture  of  one  part  sodium 
acetate  and  four  parts  soda  lime.  Heat  it  until  gas  begins  to 
form,  and  then  keep  the  temperature  as  constant  as  possible. 
Collect  the  gas  over  water,  and  recognize  it  as  the  same  gas 
as  in  d,  i.e.,  marsh  gas.  Dissolve  the  contents  of  the  tube 
when  cool  in  water.  Filter,  crystallize,  and  recognize  as 
sodium  carbonate.  The  calcium  oxid  in  the  soda  lime  is 
added  only  to  prevent  fusion. 

It  is  difficult  to  show  to  the  beginner  the  composition  of 
sodium  acetate.  Suffice  it  to  say  however  that  it  is  com- 
posed of  sodium,  carbon,  oxygen,  and  hydrogen. 

Hence  we  have 

Sodium 

Sodium         Sodium 

Carbon  Carbon 

+  Oxygen     =  Carbon  + 


f.  Ethylme.  In  a  1000  cc.  flask  fitted  with  a  funnel  and 
exit  tube,  place  10  cc.  of  95% 
alcohol.  Add  55  cc.  of  concen- 
trated sulfuric  acid  and  heat  gently. 
Collect  the  gas  evolved  over  water 
and  get  its  properties.  Have 
ready  a  eudiometer  tube  as  in  d, 
and  fill  it  one  fourth  full  of  the  gas 
dried  by  being  passed  through  sul- 
furic acid.  Allow  the  sparks  from 
an  induction  coil  to  pass»through 
the  gas,  and  notice  the  carbon  deposited  upon  the  platinum 


54        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

points.     Recognize  the  remaining  gas  as   hydrogen.     This 
gas  is  called  ethylene,  or  olefiant  gas. 

Since  sulfuric  acid  has  a  strong  affinity  for  water,  we  are 
justified  in  believing  that  the  action  of  the  acid  upon  the 
alcohol  was  to  abstract  hydrogen  and  oxygen  from  it  in  the 
form  of  water.  Hence  alcohol  is  composed  of  carbon,  oxy- 
gen, and  hydrogen. 

Carbon 

TT   ,  Hydrogen  ,  Carbon 

Hydrogen  =  „  J 

Oxygen         °Xygen         HydroSen 

Remark.  Carbon  and  hydrogen  unite  in  different  pro- 
portions, and  form  a  large  class  of  compounds  called  hydro- 
carbons. These  are  more  properly  studied  in  that  division 
of  chemistry  called  Organic  Chemistry. 


EXPERIMENT    34 
Nature  of  Flame 

a.  Light  a  candle,  and  press  a  glass  plate  down  upon  the 
flame.  Notice  that  combustion  is  taking  place  only  on 
the  outside  of  the  flame.  Remove  the  glass  plate,  and,  into 
the  dark  interior,  insert  the  end  of  a  small  tube  about  i  mm. 
bore.  Hold  a  lighted  match  to  the  other  end,  and  notice  the 
flame.  The  dark  interior  must  be  composed  of  unburned 
gas.  Note  the  unburned  carbon  (soot)  on  the  glass  plate. 
Does  white-hot  carbon  give  light  ?  To  what  is  the  luminosity 
of  the  flame  due? 

-•"'  b.  Examine  the  flame  of  a  Bunsen  burner  in  the  same 
way  when  the  air  holes  are  shut.  Draw  a  diagram  of  the 
Bunsen  flame  when  the  air  holes  are  open. 

c.  Hold  an  inverted  test  tube  over  a  candle  or  Bunsen 
flame  for  a  few  moments.  Remove  the  tube,  and  test  the 


ELEMENTS  AND  COMPOUNDS  55 

gas  in  it  with  lime  water,  proving  the  presence  of  carbon 
dioxid.  Hold  a  cold  glass  plate  over  a  flame  for  a  moment, 
and  note  the  condensed  water.  Hence  two  products  of  the 
burning  are  carbon  dioxid  and  water. 

d.  Bring  down  a  piece  of  wire  gauze  upon  the  flame  of  a 
Bunsen  burner.     Note    that    there    is    no  flame   above    the 
gauze.     Now  hold  a  lighted    match  over  the  gauze.     Was 
unburned  gas  passing  through  the  gauze  ?     Now  turn  off  the 
gas  altogether.     Hold  the  gauze  about  two  inches  above  the 
burner.     After   turning    on   the  gas  again,   hold    a   lighted 
match  above  the  gauze.     The  gas  will  now  burn  above  the 
gauze,  but   not  below   it.*     Perform   the  first  part  of  this 
experiment  a  second  time,  and  hold  the  gauze  on  the  flame 
for  some  time.     After  a  while  the  gauze  becomes  hot,  and  the 
gas  above  it  takes  fire. 

e.  Have  ready  a  platform  balance.     Cut  a  piece  of  wire 
netting  ten  inches  long  by  five  inches  wide.     Bend  the  net- 
ting into  the  form  of  a  cylinder,  so  that  its  longer  side  is  the 
altitude  of  a  cylinder.      Divide  the  cylinder  into  two  equal 
parts  by  means  of  a  circular  piece  of  netting  fastened  inside 
the  cylinder  by  means  of  wire.     Fill  the   upper  half  about 
one  quarter  full  of  broken  pieces  of  sodium  hydroxid.    Then 
place  upon  this  a  second  circular  piece  of    netting,  and,  in 
the    compartment    thus    made,   place   an   equal    quantity  of 
fused  calcium   chlorid.     Place    the   cylinder  upon  one  plat- 
form of  the  balance  with  the  empty  half  over  a  candle  about 

*  The  miners  safety  lamp  depends  upon  this  principle.  This  lamp 
is  simply  an  ordinary  oil  lamp  having  the  flame  surrounded  by  wire 
gauze.  The  explosive  mixtures  of  gases  (air  and  compounds  of  carbon 
and  hydrogen),  which  occur  in  mines,  may  pass  through  the  gauze  and 
burn  inside  the  lamp  ;  but  outside,  on  account  of  the  relatively  high 
"  kindling  temperatures "  of  carbon  and  hydrogen,  they  do  not  take 
fire. 


56        AN     ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

three  quarters  of  an  inch  long,  and,  upon  the  other,  place 
weights  enough  to  balance  it  exactly.  Light  the  candle,  and 
watch  the  effect  upon  the  equilibrium.  Be  careful  not  to 
place  the  apparatus  in  a  draught.  What  is  the  relation  be- 
tween the  weight  of  the  products  formed  by  the  burning 
candle  and  the  weight  lost  by  the  candle  ? 


EXPERIMENT    35 
Hard  and  Soft  Water 

a.  Permanently  Hard  Water.     Place  about  5  grms.  of  cal- 
cium sulfate  in  50  cc.  of  distilled  water  in  a  beaker ;  stir,  and 
filter  it.     Place  10  cc.  of  this  water  in  a  test  tube,  and  add 
from  a  soap  solution,*  drop  by  drop,  until  a  permanent  froth 
is  obtained,  counting  the  number  of  drops  necessary.     In  the 
same  way,  test  10  cc.  of  distilled  water.     Now  boil  the  re- 
maining 40  cc.  in  the  beaker,  and,  when  it  is  cool,  test  10  cc. 
of  it  again  with  soap  solution.     Does  boiling  make  the  water 
soft? 

b.  Temporarily  Hard   Water.     Into  a   beaker  containing 
50  cc.  of  lime  water,  pass  carbon  dioxid  until  the  precipitate 
first  formed  is  dissolved.      Test  10  cc.  of  this  with  soap  so- 
lution the  same  as  in  a.      Boil  the  remaining  40  cc.  until  the 
precipitate  all  returns.      Then  filter  and  test  10  cc.  of  it  with 
soap  solution.     Has  the  water  become  soft  ? 

Remark.  Water  containing  calcium  sulfate  in  solution  is 
called  permanently  hard  water.  It  cannot  be  made  soft  by 
boiling,  since  calcium  sulfate  is  about  as  soluble  in  hot  water 
as  in  cold. 

*  To  make  a  soap  solution,  place  I  grm.  of  castile  soap  shavings  in 
10  cc.  of  distilled  water,  and  dissolve  as  much  as  possible.  Then  pour 
off  the  clear  solution. 


ELEMENTS  AND  COMPOUNDS  57 

Water  containing  calcium  carbonate  held  in  solution  by 
carbonic  acid  is  called  temporarily  hard  water,  since  it  can  be 
made  soft  by  boiling,  thus  driving  off  the  carbon  dioxid  and 
precipitating  the  calcium  carbonate. 

EXPERIMENT    36 
Nitrogen 

Float  on  water  in  a  pneumatic  trough  a  flat  cork  fitted 
with  the  bowl  of  a  deflagrating  spoon  containing  a  piece  of 
phosphorus  about  the  size  of  a  pea. 
Ignite  the  phosphorus,  and  hold  a 
bell  jar  over  it.  Place  a  glass  plate  / 
under  the  jar,  when  it  is  cool ;  keep-  | 
ing  the  water  inside  and  outside  at 
the  same  level ;  remove  the  jar,  and 
shake  it.  Note  that  one  fifth  of  the 
air  was  oxygen.  Get  the  properties 
of  the  remaining  four-fifths  of  the  air. 
This  gas  is  called  nitrogen. 

Remark  i.  Air  is  a  mechanical  mixture  and  not  a  chemi- 
cal compound.  The  following  are  the  best  proofs  of  this 
statement :  — 

1 .  Nitrogen  and  oxygen  are  not  present  in  the  proportions 
in  which  they  would  be,  if  they  were  chemically  united. 

2.  The  composition  of  air  is  slightly  variable.     (See  Ex- 
periment 2,  Part  II.) 

3.  A  mixture  of  nitrogen  and  oxygen  in  the  proportion  of 
4  to  i  has  the  same  properties  as  air. 

4.  Air  is  slightly  soluble  in  water.     But  air  dissolved  in 
water  does  not  contain  the  same  proportions  of  nitrogen  and 
oxygen  as  atmospheric  air.     The  water  takes  up  the  two 
gases  according  to  their  respective  solubilities. 


58         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Remark  2.  Since  organized  bodies,  both  vegetable  and 
animal,  all  have  as  their  principal  constituent  the  element 
carbon,  the  oxidation  of  these  bodies,  either  by  the  rapid 
process  of  combustion  or  the  slower  one  of  decay,  must  be 
accompanied  by  an  evolution  of  carbon  dioxid.  From  these 
sources  the  carbon  dioxid  in  the  atmosphere  is  obtained. 
When  the  life  process,  whatever  it  may  be,  ceases  to  act,  the 
oxygen  of  the  air,  aided  by  micro-organisms  known  under  the 
different  names  of  germs,  microbes,  bacilli,  etc.,  immediately 
attacks  the  compounds  of  which  the  body  is  composed,  and, 
by  more  or  less  complicated  processes,  changes  them  into 
other  compounds,  one  of  which  is  carbon  dioxid.  Even  dur- 
ing life  the  process  goes  on,  and  sometimes  it  is  absolutely 
necessary  to  the  living  organism,  as  in  the  case  of  the  oxi- 
dation that  is  taking  place  continually  in  the  bodies  of  ani- 
mals. The  oxygen  of  the  air  is  taken  into  the  lungs,  where 
it  is  taken  up  by  the  blood.  A  certain  substance  called 
haemoglobin,  which  is  found  in  the  red  blood  corpuscles, 
unites  with  the  oxygen,  and  thus  it  is  carried  to  the  remotest 
part  of  the  body.  The  haemoglobin  gives  up  its  oxygen 
wherever  there  is  material  whose  oxidation  the  economy  of 
the  life  process  requires.  The  products  of  this  oxidation  are 
carried  back  to  the  lungs,  where  they  are  expelled  into  the 
air.  The  chemical  energy,  stored  in  the  compounds  that 
are  thus  oxidized,  is  tranformed  into  heat,  and,  in  this  man- 
ner, keeps  up  the  necessary  warmth  of  the  body.  The  car- 
bon dioxid  that  is  expelled  into  the  air  from  the  lungs  of 
living  animals  would  soon  increase  the  quantity  of  that  gas 
in  the  atmosphere,  were  it  not  for  the  reciprocal  action  of 
plants.  Just  as  animals  require  oxygen,  so  plants  require 
carbon  dioxid.  The  leaves  of  the  plant  admit  the  carbon 
dioxid  through  small  apertures  on  their  under  sides.  In  the 
leaves  there  is  a  green  substance  called  chlorophyll.  This, 


ELEMENTS  AND  COMPOUNDS  59 

acted  upon  by  sunlight,  decomposes  the  carbon  dioxid ;  and 
the  carbon,  together  with  hydrogen,  oxygen,  and  other  ele- 
ments that  have  been  obtained  through  the  roots  in  the  form 
of  water  and  other  compounds,  is  deposited  throughout  the 
plant  structure.  The  oxygen  that  is  left  from  the  carbon 
dioxid  is  expelled  into  the  air.  In  this  manner  the  relative 
quantities  of  oxygen  and  carbon  dioxid  in  the  air  are  kept 
practically  constant. 

It  is  well  to  note  here  the  difference  between  organic  bod- 
ies and  organized  bodies.  The  life  process  shapes  the  organic 
matter  into  structures  called  cells,  which,  taken  together,  make 
up  the  organized  body.  The  study  of  the  chemistry  of  plant 
and  animal  life  belongs  to  a  special  branch  of  chemistry, 
called  Physiological  Chemistry. 


EXPERIMENT    37 
Nitric  Acid 

a.  Preparation.     In  a  glass-stoppered  retort, place  30  grms. 
of  powdered  sodium  nitrate  (Chili  saltpeter).     Add  10  cc.  of 
concentrated  sulfuric  acid,  being  careful  to  allow  none  of  it 
to  fall  into  the  neck  of  the  retort.     Have   the  neck  of  the 
retort  extending  into  a  100  cc.  flask  over  which  cold  water  is 
kept  running.     Heat  the  retort  gently,  and  distill  over  the 
liquid.     Get  its  properties.     Dilute  a  little  with  water,  and 
test  with  blue  litmus.     This  is  called  nitric  acid.     When  the 
residue  is  cool,  dilute  it  in  the  retort  with  water,  evaporate, 
crystallize,  and  recognize  it  as  sodium  sulfate. 

b.  i .    Analysis  of  Nitric  Acid.     Pour  into  a   i  oo  cc.  flask, 
fitted  with  a  one-holed  cork  stopper  and  delivery  tube,  a  little 
dilute  nitric  acid,  and  add  a  few  pieces  of  magnesium  rib- 
bon.    Catch  the  gas  evolved,  and  recognize  it  as  hydrogen. 


6O        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Since  the  hydrogen  could  not  have  come  from  the  magne- 
sium, or  from  the  magnesium  and  water,  it  must  have  come 
from  the  nitric  acid.  Hence  one  constituent  of  nitric  acid 
is  hydrogen  (see  def.  acid,  page  29). 


2.  Have  ready  an  apparatus  for  obtaining  a  steady  stream 
of  dry  carbon  dioxid  (see  Experiment  26).  Connect  this 
with  a  test  tube  fitted  with  an  exit  tube  and  an  entrance 
tube  extending  almost  to  the  bottom.  Place  about  4  cc.  of 
concentrated  nitric  acid  in  the  test  tube.  Connect  the  test 
tube  with  a  combustion  tube  containing  clean  copper  filings. 
Have  ready  a  glass  vessel  containing  a  moderately  strong 
sodium  hydroxid  solution.  Pass  carbon  dioxid  through  the 
apparatus  to  remove  the  air  ;  then  gently  heat  the  nitric  acid, 
having  meanwhile  heated  the  copper  filings  to  redness.  Col- 
lect the  gas  evolved  over  the  sodium  hydrate  solution,  and 


ELEMENTS  AND  COMPOUNDS 


6l 


recognize  it  as  nitrogen.  The  sodium  hydroxid  absorbs  the 
carbon  dioxid,  which  is  simply  used  as  a  carrier  for  the  nitric 
acid  fumes.  Note  the  formation  of  copper  oxid  in  the  com- 
bustion tube. 

Since  copper  does  not  affect  carbon  dioxid,  the  oxygen 
must  have  come  from  the  nitric  acid.  Therefore  nitric  acid 
contains  nitrogen  and  oxygen.  We  then  have  as  the  com- 
ponent parts  of  nitric  acid,  hydrogen,  nitrogen,  and  oxygen. 


Argument.  In  the  making  of  nitric  acid  we  started  with 
the  sodium  nitrate,  a  substance  whose  composition  we  do 
not  know.  This  acted  with  sulfuric  acid,  and  a  compound 
containing  sodium,  sulfur,  and  oxygen,  i.e.,  sodium  sulfate, 
was  formed,  together  with  another  compound  containing  hy- 
drogen, nitrogen,  and  oxygen.  We  then  have  — 

/  \       Hydrogen       Sodium      Hydrogen 

\  )  +  Sulfur       =  Sulfur   +  Nitrogen 

Sodium  nitrate      Oxygen          Oxygen      Oxygen 


62         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Knowing  that  the  tendency  of  metals  is  to  replace  the  hy- 
drogen of  acids,  we  readily  see  that  the  action  was 

Sodium          Hydrogen       Sodium        Hydrogen 
Nitrogen   +  Sulfur       =   Sulfur     -f-  Nitrogen 
Oxygen  Oxygen  Oxygen        Oxygen 

If  this  be  true,  we  ought  to  obtain  sodium  nitrate  by  neu- 
tralizing nitric  acid  with  sodium  hydroxid.  (See  Experiment 

38.) 

Remark.  Nitric  acid,  when  left  exposed  to  the  air,  gives 
off  colorless  fumes.  When  heated  or  exposed  to  sunlight, 
it  decomposes  slowly  into  water,  oxygen,  and  various  oxids 
of  nitrogen.  This  is  the  reason  the  brown  fumes  are  seen 
above  the  liquid  in  nitric  acid  bottles.  Nitric  acid  is  a  very 
valuable  oxidizing  agent  on  account  of  the  ease  with  which 
it  gives  up  its  oxygen.  The  compounds  of  nitric  acid  with 
organic  substances  are  especially  unstable. 


EXPERIMENT   38 

Neutralization  of   Nitric  Acid 

Neutralize    nitric    acid    with    sodium    and  potassium   hy- 
droxids. 

Sodium  Hydrogen       Sodium 

(1)  Oxygen     +    Nitrogen   =  Nitrogen    +      y 
Hydrogen        Oxygen          Oxygen 

Potassium       Hydrogen      Potassium 

(2)  Oxygen      +   Nitrogen  =  Nitrogen    +o 
Hydrogen         Oxygen          Oxygen  Xy*  n 


ELEMENTS  AND  COMPOUNDS  63 

EXPERIMENT    39 

Nitric  Oxid 

a.  Preparation.  Place  about  50  grms.  of  sheet  copper 
clippings  or  else  copper  turnings  in  a  250  cc.  flask  fitted 
with  funnel  and  exit  tubes.  Pour  in  through  the  funnel  tube 
just  enough  nitric  acid,  diluted  with  an  equal  volume  of  water, 
to  cover  the  copper.  Collect  the  colorless  gas  in  jars  over 
water,  and  note  its  properties.  Open  a  jar  in  the  air,  and 
note  the  properties  of  the  new  brown  gas  thus  formed.  Ignite 
a  piece  of  yellow  phosphorus  about  the  size  of  a  bean  in  a 
deflagrating  spoon,  and  hang  it  in  a  second  jar  of  the  gas  as 
in  Experiment  7  b.  When  the  burning  has  ceased  and  the 
jar  is  cool,  and  after  the  phosphoric  oxid  has  settled,  rec- 
ognize the  remaining  gas  as  nitrogen  (see  Experiment  36). 
Therefore  the  two  gases,  the  colorless  and  the  brown,  must 
be  oxids  of  nitrogen,  the  latter  containing  the  more  oxygen. 

Here  we  have  a  case  where  apparently  a  metal  does  not 
replace  the  hydrogen  of  the  acid.  The  fact  of  the  matter  is 
that  elements  at  the  moment  they  are  released  from  their 
compounds  will  unite  with  other  elements,  or  decompose 
compounds  that  they  would  not  affect  at  other  times.  In 
this  case,  the  hydrogen  we  should  naturally  expect  to  be 
evolved  attacks  the  extra  nitric  acid  instead,  forming  water 
and  nitric  oxid.  We  may  express  the  two  actions  thus  :  - 

Hydrogen      Copper 

(1)  Copper  +  Nitrogen  =  Nitrogen  +  Hydrogen 

Oxygen         Oxygen 

Hydrogen 

(2)  Hydrogen  +  Nitrogen  =  *ydr°gen  +  ™r°gen 

Oxygen          Oxygen         Oxygen 


64          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Or  we  might  write  the  two  together  thus  :  — 

Hydrogen      Copper 

*  Hydrogen     Nitrogen 

Copper  +  Nitrogen  =  Nitrogen  +     *      6     +.        6 
T_  Oxygen         Oxygen 

Oxygen          Oxygen 

b.  Sulfuric  Acid  by  means  of  Nitric  Ox  id.  Have  ready  a 
large  flask  (capacity  about  3  or  4  liters),  loosely  fitted  with 
a  five-holed  stopper.  Let  four  of  the  holes  contain  entrance 
tubes,  bent  at  right  angles,  and  extending  to  the  bottom  of 


the  flask.  Let  the  fifth  hole  contain  an  exit  tube.  Fit  three 
250  cc.  flasks  with  funnel  tubes  and  exit  tubes,  and  arrange 
them  so  that  they  may  be  connected  with  three  of  the  en- 
trance tubes  of  the  large  flask.  Fill  the  first  of  the  small 


ELEMENTS    AND    COMPOUNDS  65 

flasks  about  half  full  of  water ;  in  the  second,  place  copper 
and  sulfuric  acid  (as  in  Experiment  9  g) ;  and  in  the  third, 
place  copper  and  nitric  acid  (as  in  a  of  this  experiment). 
Connect  the  fourth  entrance  tube  with  a  pair  of  bellows. 
Allow  steam,  air,  nitric  oxid,  and  sulfur  dioxid  to  enter  the 
large  flask.  Note  the  decolorization  of  the  brown  gas  that 
was  formed  when  the  nitric  oxid  first  entered.  Evidently 
the  oxygen  that  the  nitric  oxid  took  on  has  been  taken  away 
by  some  other  compound.  Steam  does  not  have  that  prop- 
erty ;  hence  it  must  have  been  the  sulfurous  acid  formed  by 
the  sulfur  dioxid  and  water  (see  Experiment  9  d).  But  if 
sulfurous  acid  takes  on  oxygen,  it  becomes  sulfuric  acid  (see 
Experiment  9,  */and/). 

After  the  operation  has  been  going  on  for  some  little  time 
in  the  large  flask,  remove  the  liquid  that  has  collected,  and 
recognize  it  as  sulfuric  acid. 

This  experiment  illustrates  the  principle  that  is  used  in 
the  manufacture  of  sulfuric  acid  on  a  large  scale.  (See 
page  173)- 


EXPERIMENT    40 

Elements  in  Nascent  State 

Make  a  solution  of  a  few  crystals  of  potassium  perman- 
ganate. Through  this,  pass  a  stream  of  hydrogen  from  a 
hydrogen  generator.  Is  there  any  change  in  color  ?  Then 
remove  the  generator,  and  add  a  little  concentrated  sulfuric 
acid.  Is  there  now  any  change  in  color  ?  Add  to  this  a 
few  pieces  of  zinc.  What  results  ? 

Here  the  hydrogen  that  first  passed  through  the  solution 
evidently  had  no  effect.  In  the  second  case,  where  the 
hydrogen  was  set  free  by  the  zinc  in  the  presence  of  the  per- 


66        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

manganate,  it  evidently  decomposed  it.  Elements  at  the  in- 
stant they  leave  their  compounds  have  a  stronger  chemical 
attraction  for  other  elements,  and  are  said  to  be  in  the 
nascent  state. 

EXPERIMENT    41 
Ammonia 

a.  Ammonia  from  Nitric  Oxid  and  Hydrogen.  Fill  a  gas 
holder  with  two  volumes  of  nitric  oxid  and  five  volumes  of 
hydrogen,  made  respectively  from  copper  and  nitric  acid, 
and  zinc  and  sulfuric  acid.  Place  loosely  in  a  glass  com- 


bustion tube  about  2  grms.  of  platinized  asbestos,  and  con- 
nect the  tube  with  the  gas  holder  in  such  away  that  the 
mixed  gases  must  pass  through  sulfuric  acid  in  a  catch  bottle. 
To  the  other  end  of  the  tube,  attach  a  piece  of  rubber  tubing 
about  two  feet  long.  Allow  the  mixed  gases  to  pass  slowly 
through  the  apparatus  until  the  air  is  expelled;  then,  when 
no  brown  color  is  seen,  heat  the  platinized  asbestos  with  a 


ELEMENTS  AND  COMPOUNDS  6? 

Bunsen  flame.  Notice  the  water  that  collects  on  the  cooler 
part  of  the  tube.  Get  the  odor  of  the  gas  that  escapes,  and 
test  the  gas  with  moist  red  litmus  paper. 

Since,  in  the  mixed  gases,  there  are  only  the  elements 
hydrogen,  nitrogen,  and  oxygen,  and  since  water  is  formed, 
the  escaping  gas  must  be  either  nitrogen,  hydrogen,  oxygen, 
a  compound  of  nitrogen  and  oxygen,  or  else  a  compound  of 
nitrogen  and  hydrogen.  On  account  of  its  properties,  it  can- 
not be  any  of  the  first  four  (see  Experiments  36,  15  a,  6, 
and  39),  so  it  must  be  composed  of  nitrogen  and  hydrogen. 
We  call  the  gas  ammonia. 

Nitrogen  Nitrogen        Hydrogen 

Oxygen          HydroSen  =  Hydrogen  +  Oxygen 

b.  Ammonia  from  Organic  Matter  containing  Nitrogen. 
Make  a  mixture  of  freshly  slacked  lime  and  dry  sodium 
hydroxid.  Cut  up  a  piece  of  flannel  or  horn  into  small  bits, 
and  heat  the  pieces  in  a  test  tube  with  the  above  mixture. 
Recognize  the  gas  that  escapes  as  ammonia. 


EXPERIMENT    42 
Ammonium  Chlorid 

Allow  ammonia  gas  and  hydrochloric  acid  gas  to  come  in 
contact  with  each  other,  and  note  the  white  fumes  that  con- 
dense in  solid  form.  Taste  the  new  substance. 

This  new  compound  must  be  composed  of  nitrogen,  hy- 
drogen, and  chlorin ;  but  there  must  be  more  hydrogen  in  it 
than  in  ammonia.  To  indicate  this  we  name  the  substance 
ammoniww  chlorid. 

Nitrogen 

S1?*"1  +  ?£     ^  =  Hydrogen 
Hydrogen     Chlorin 


68         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

EXPERIMENT    43 

Ammonia  from  Ammonium  Chlorid 

Mix  intimately  20  grms.  of  ammonium  chlorid  and  40 
grms.  of  calcium  oxid.  Put  the  mixture  into  a  copper  re- 
tort,* or  into  a  large  piece  of  a  gas  pipe  closed  with  a  plug 
at  one  end.  Connect  this  with  a  tube  containing  calcium 
chlorid  to  dry  the  evolved  gas,  and  after  heating  recognize 
the  gas  as  ammonia.  Note  that  the  calcium  chlorid  absorbs 
water.  Collect  several  jars  full  by  displacement  of  air,  i.e., 
allow  a  delivery  tube  to  reach  to  the  bottom  of  the  jar,  but 
have  the  jar  inverted,  since  ammonia  is  lighter  than  air. 
Open  a  jar  under  water,  and  note  the  great  solubility  of  the 
gas. 

Since  water  and  ammonia  are  formed  by  the  chemical 
action,  the  only  elements  left  to  account  for  are  the  calcium 
and  the  chlorin.  These  must  have  united  to  form  calcium 
chlorid,  which  action  may  be  verified  by  dissolving  the  con- 
tents of  the  retort,  filtering  and  crystallizing.  Therefore  we 
have 

Calcium  _  Nitrogen        Hydrogen      Calcium 

7.    .  Oxygen   ~~  Hydrogen       Oxygen          Chlorin 

Cnlorin 

Remark.  When  ammonia  dissolves  in  water,  we  call  the 
product  ammonium  hydroxid,  that  is, 

Nitrogen 
Hydrogen 
Oxygen 
Hydrogen 

*  If  the  chlorid  and  lime  are  dry,  the  experiment  may  be  performed 
in  an  ordinary  flask. 


ELEMENTS  AND  COMPOUNDS  69 

The  two  elements  nitrogen  and  hydrogen,  as  they  appear 
in  ammonium  chlorid  and  ammonium  hydroxid,  are  capable 
of  passing  from  compound  to  compound  without  separating, 
and  act  just  as  if  they  were  a  metal.  Such  a  combination  of 
elements  is  called  a  compound  radical. 

EXPERIMENT    44 
Ammonium  Amalgam 

To  show  how  much  like  a  metal  the  radical  ammonium 
acts,  it  is  possible  to  make  an  ammonium  amalgam  with 
mercury.  Into  a  strong  solution  of  ammonium  chlorid,  drop 
a  piece  of  sodium  amalgam  (see  Experiment  io*/).  As 
the  ammonium  amalgam  decomposes,  notice  the  fumes  of 
ammonia. 

EXPERIMENT    45 
Neutralization  of  Acids  with  Ammonium  Hydroxid 

Neutralize  dilute  hydrochloric,  sulfuric,  and  nitric  acids 
with  ammonium  hydroxid,  and  verify  the  following  :  — 

Nitrogen 

Hydrogen       Hydrogen  Hydrogen 


(  =  + 

Oxygen      T  Chlorin  Cn\0rin          Oxygen 

Hydrogen 


(} 


Nitrogen 
Hydrogen 
Oxygen 
Hydrogen 

Hydrogen 
Sulfur 
Oxygen 

Nitrogen 
Hydrogen 
=  Sulfur 
Oxygen 

Hydrogen 
Oxygen 

Nitrogen 

H  dro  en 

Nitrogen 

Hydrogen 
Oxygen 
Hydrogen 

Nitrogen 
Oxygen 

Hydrogen 
~~  Nitrogen 
Oxygen 

Hydrogen 
Oxygen 

7O        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Evaporate  the  solution.  Place  a  little  of  each  of  the  salts 
formed  in  small  tubes  closed  at  one  end,  and  heat.  The  phe- 
nomenon in  the  case  of  the  first  two  is  called  sublimation. 
Name  the  salts. 


EXPERIMENT    46 
Nitrous  Oxid 

In  a  stoppered  glass  retort,  place  15  grms.  of  ammonium 
nitrate.  Connect  the  retort  with  three  catch  bottles,  the 
first  empty,  the  second  containing  a  dilute  solution  of  sodium 
hydroxid,  and  the  third  containing  a  solution  of  iron  sulfate. 
Heat  the  retort  gently.  Note  the  formation  of  water  in  the 


first  catch  bottle.  The  sodium  hydroxid  will  absorb  any 
chlorin  which  may  be  obtained  from  the  commercial  am- 
monium nitrate,  and  the  iron  sulfate  will  absorb  any  nitric 
oxid  which  will  be  formed  if  the  ammonium  nitrate  is  heated 
too  hot.  Collect  the  gas  over  warm  water,  since  it  is  soluble 
in  cold.  Note  its  properties,  especially  odor,  taste,  and 
relation  to  combustion. 


ELEMENTS    AND    COMPOUNDS  7 I 

Since  ammonium  nitrate  is  composed  of  nitrogen,  hydro- 
gen, and  oxygen,  and  since  water  is  formed  in  the  experi- 
ment, the  new  gas  must  be  either  nitrogen,  a  compound  of 
nitrogen  and  hydrogen,  or  a  compound  of  nitrogen  and 
oxygen.  On  account  of  the  properties  of  the  first  and  sec- 
ond, and  on  account  of  its  supporting  combustion,  it  must 
be  the  third.  Here  we  have  another  oxid  of  nitrogen,  which 
is  called  nitrous  oxid  to  distinguish  it  from  nitric  oxid  and 
nitric  peroxid. 

Nitrogen 

Hydrogen  _  Hydrogen       Nitrogen 

Nitrogen   ~  Oxygen  Oxygen 

Oxygen 

Remark.  This  oxid  of  nitrogen  is  commonly  called  laugh- 
ing gas.  When  taken  into  the  lungs,  it  acts  as  an  anaes- 
thetic, and  is  hence  valuable  in  certain  surgical  operations. 
For  this  reason  it  is  used  by  dentists.  Its  comparative  in- 
stability is  shown  by  the  fact  that  it  gives  up  its  oxygen 
easily,  thus  supporting  combustion.  Whatever  the  forces 
are  that  hold  elements  together,  they  are  evidently  in  this 
case  in  "  unstable  equilibrium"  i.e.,  if  a  substance  easily 
oxidized  is  brought  in  contact  with  nitrous  oxid  under  proper 
conditions,  the  oxid  gives  up  its  oxygen.  In  the  case  of 
nitric  oxid,  however,  the  equilibrium  seems  to  be  more  stable  ; 
and,  although  there  is  more  oxygen  in  the  compound,  it  is 
held  there  more  firmly.  From  Experiment  39,  we  learn  that 
nitric  oxid  will  unite  with  more  oxygen,  thus  giving  a  still 
higher  oxid.  It  is  therefore  a  reducing  agent.  On  the 
other  hand,  the  oxygen  thus  obtained  by  the  nitric  oxid  is 
not  held  firmly,  and  the  peroxid  formed  will  give  it  up  read- 
ily. Nitrogen  peroxid  is  therefore  an  oxidizing  agent. 


/2         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

EXPERIMENT    47 

Analysis  of  Nitrous  Oxid 

Collect  5  cc.  of  the  gas  over  mercury  in  a  eudiometer  tube 
with  platinum  wires  fused  in  the  closed  end.  Add  to  this 
an  equal  volume  of  hydrogen.  Let  both  gases  be  dry.  Pass 
sparks  through  the  mixture  by  means  of  an  induction  coil. 
Recognize  the  remaining  gas  as  nitrogen. 

Nitrogen  ,  Hydrogen 

~  +  Hydrogen  =  Nitrogen  +  ^J 

Oxygen  Oxygen 

EXPERIMENT   48 
Arsenic 

Perform  all  experiments  with  arsenic  under  a  good  hood. 
Be  careful  not  to  breathe  any  arsenic  fumes,  since  they  are 
deadly  poisons. 

a.  Properties.     Take  a  bit  of  the  element  arsenic,  and  note 
as  many  properties  as  you  can.     Place  a  small  piece  about 
2  mm.  in  diameter  in  a  glass  tube  closed  at  one   end.     Heat 
the  tube  in  the   Bunsen  flame,  noticing  that   arsenic  passes 
immediately  from  the  solid  to  the  gaseous  state,  after  which 
it  again  condenses  on  the   cooler  part   of  the  tube,  forming 
the  so-called  arsenic  mirror. 

b.  Arsenic  Oxid.     Oxidize  a  bit  of  arsenic  about  the  size 
of  a  pin  head  under  a  hood.     Be  careful  not  to  breathe  the 
fumes.     You  can  hardly  avoid  noticing  the  garlic-like  odor 
of  arsenic. 

Arsenic 
Arsenic  +  Oxygen  = 


ELEMENTS    AND    COMPOUNDS 


73 


c.  Reduction  of  Arsenic  Oxid  by  means  of  Carbon.     Mix  a 
little  arsenic  oxid  with  powdered  charcoal,  and  place  it  in  a 
tube  like  that  used  in  a.     Heat  the  tube  and  verify  the  fol- 
lowing :  — 

Arsenic    ,  Carbon    , 

_  f  Carbon  =  _.  +  Arsenic 

Oxygen  Oxygen 

d.  Arsin  (Hydrogen  Arsenid).     Connect  a  flask  for  gen- 
erating hydrogen  (zinc  and  sulfuric  acid)  with  a  U  tube  con- 


taining calcium  chlorid,  and  then  join  this  with  a  glass  tube 
narrowed  in  several  places  and  finally  drawn  to  a  small  open- 
ing. When  the  air  is  expelled  from  the  apparatus,  light  the 
hydrogen  jet  and  hold  a  clean  crucible  cover  in  the  flame 
for  a  moment.  Then  add  a  little  arsenic  or  arsenic  oxid 
through  the  funnel  tube  of  the  hydrogen  generator.  Note 
the  change  of  color  in  the  flame  and  the  vapors  of  arsenic 


74         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

oxid  that  arise.*  Now  again  place  the  porcelain  crucible 
cover  in  the  flame,  and  note  the  arsenic  mirror  formed. 
Apply  the  flame  of  a  Bunsen  burner  to  one  of  the  bulbs  of 
the  tube,  and  note  the  formation  of  the  mirror  in  the  narrow 
part,  just  beyond  the  heated  portion.  Test  the  mirror  on  the 
crucible  cover  with  a  solution  of  bleaching  powder. 

We  have  seen  before  that  elements,  when  in  the  nascent 
state,  combine  more  easily  with  other  elements.  Here  we 
have  nascent  hydrogen  uniting  with  arsenic,  forming  hydro- 
gen arsenid,  or  arsin. 

Arsenic 
Arsenic  +  Hydrogen  =  TT    , 

Hydrogen 

Remark.  There  are  two  acids  of  arsenic,  similar  in  their 
chemical  properties  to  phosphorous  and  phosphoric  acids. 
They  are  called  arsenious  and  arsenic  acids.  Both  are  com- 
posed of  hydrogen,  arsenic,  and  oxygen,  the  latter  contain- 
ing the  greater  amount  of  oxygen. 


EXPERIMENT    49 
Antimony 

a.  Properties.  Examine  a  piece  of  antimony,  and  note  its 
properties.  See  whether  you  can  obtain  a  mirror,  as  with 
arsenic. 

*  In  this  experiment  we  have  both  complete  and  incomplete  com- 
bustion of  arsin.  In  the  first  case,  the  products  of  combustion  are 
evidently  water  and  arsenic  oxid  ;  in  the  second  they  are  water,  ar- 
senic oxid,  and  arsenic.  In  the  same  way,  when  carbon  bisulfid  burns,  we 
obtain  either  carbon  dioxid  and  sulfur  dioxid,  or  else  carbon  dioxid,  sul- 
fur dioxid,  and  sulfur ;  when  hydrogen  sulfid  burns,  we  obtain  either 
water  and  sulfur  dioxid,  or  else  water,  sulfur  dioxid,  and  sulfur ;  and 
when  coal  gas  burns,  we  obtain  either  water  and  carbon  dioxid,  or  else 
water,  carbon  dioxid,  and  carbon.  It  is  evident,  the-efore,  that  the  affin- 
ity of  oxygen  for  these  elements  in  combination  is  not  so  great  as  when 
they  are  free. 


ELEMENTS    AND    COMPOUNDS  75 

b.  Antimony  Oxid.  Oxidize  a  small  piece  of  antimony, 
and  note  the  color  of  the  oxid.  Drop  a  bit  of  melted  anti- 
mony on  a  sheet  of  paper. 

,   n  Antimony 

Antimony  +  Oxygen  = 


c.  Antimony  Chlorid.     Into  a  jar  of  chlorin,  drop  a  little 
powdered  antimony,   and  note    the    formation  of   antimony 
chlorid. 

Antimony  +  Chlorin  =  _ 

Chlorin 

d.  Stibin  (Hydrogen  Antimonid).     Perform  an  experiment 
exactly  like   Experiment  48  d,  with  the  exception  of    using 
antimony  or   antimony   oxid   instead   of  arsenic.     Test  the 
mirror  with  a  solution  of  bleaching  powder. 

Antimony 
Antimony  +  Hydrogen  =  _   , 

Hydrogen 

EXPERIMENT   50 
Bismuth 

a.  Properties.     Take    a    piece    of   bismuth,  and    note   its 
properties. 

b.  Bismuth  Oxid.     Oxidize  a  small   piece  of  bismuth  by 
heating  it  with  the  blowpipe  on  charcoal.     Note  the  color  of 

the  oxid. 

_  Bismuth 

Bismuth  +  Oxygen  =  _ 

Oxygen 

c.  Bismuth  Nitrate.     Heat  a  small  piece  of  bismuth  to- 
gether with  a  little  nitric  acid.     The   action   of  nitric   acid 
upon  bismuth  is  similar  to  its  action  upon  copper  (see  Ex- 
periment 39). 

Hydrogen       Bismuth 

Bumuth  +  N,trogen   =  NUrogen  +    ^    *     +  Oxygen 
Oxygen          Oxygen 


/6         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

d.  Basic  Bismuth  Nitrate.     To  the   solution  formed  in  <r, 
add  water   and   note  the    formation  of  a  white  precipitate. 
This  is  called  basic  bismuth  nitrate.     (See  Remark  under  Ex- 
periment 1 8.) 

e.  Bismuth  Sulfid.     Dilute  a  solution  of  bismuth  nitrate ; 
then  add  just  enough  nitric  acid  to   prevent  the  formation  of 
the  basic  nitrate,  and  pass  through  it  a  stream  of  hydrogen 
sulfid  made  as  in  Experiment  32  c.     Note  the  color  of  the  pre- 
cipitate formed. 

Bismuth  Hydrogen 

Nitrogen  +     ^J       =  Sul£ur      +  Nitrogen 
Oxygen  Oxygen 


EXPERIMENT    51 
Cadmium 

a.  Properties.     Examine  a  piece  of  cadmium,  and  note  its 
properties. 

b.  Cadmium  Oxid.    Oxidize  cadmium  on  the  charcoal,  and 
note  the  color  of  the  oxid. 

Cadmium 
Cadmium  +  Oxygen  =  Qxygen 

c.  Action  of  Acids  upon  Cadmium.     Try   the  effect   of  hy- 
drochloric, sulfuric,  and  nitric  acids  on  cadmium. 

Hydrogen       Cadmium 
(,)    Cadmium  +  =  h  Hydrogen 


Hydrogen      Cadmium 

(2)  Cadmium  +  Sulfur       =  Sulfur      +  Hydrogen 

Oxygen          Oxygen 

Hydrogen      Cadmium  ^ 

(3)  Cadmium  +  Nitrogen  -  Nitrogen  +     *      *     +  Q 

Oxygen          Oxygen 


ELEMENTS  AND  COMPOUNDS  // 

d.    Cadmium    Sulfid.     Try  the    effect   of  hydrogen  sulfid 
upon  cadmium  nitrate. 

Cadmium  Hydrogen 


Nitrogen    +  gulfur        =  Sulfu +  Nitrogen 

Oxygen  Oxygen 

EXPERIMENT    52 
Mercury 

For  properties  and  oxid  see  Experiment  5. 

a.  Mercurous   and  Mercuric  Nitrates.     Try  the   effect  of 
nitric  acid  upon  a  globule  of  mercury. 

Hydrogen       Mercury 

Mercury  +  Nitrogen   =  Nitrogen  +  *y<Jr°g<m  +  Nltr°«en 

r\  Oxygen  Oxygen 

Oxygen  Oxygen 

Remark.  Mercury  unites  with  the  nitrogen  and  oxygen  of 
nitric  acid  in  two  proportions,  the  one  containing  the  smaller 
quantity  of  the  acid  radical  being  called  mercurous  nitrate, 
and  the  other  mercuric  nitrate. 

b.  Mercurous  Chlorid.     To  a  solution  of  mercurous  nitrate, 
add  hydrochloric  acid  until  no  further  precipitate  is   formed. 
The  precipitate  is  mercurous  chlorid. 

Mercury  Hydrogen 

M  Hydrogen      Mercury 

Nitrogen  +     *  =  _. .    .  *  +  Nitrogen 

_  Chlorm          Chlorm 

Oxygen  Oxygen 

EXPERIMENT    53 
Lead 

a.  Properties.     State  the  properties  of  lead. 

b.  Lead  Oxid.     Oxidize  a  small  piece  of  lead  on  charcoal, 
and  state  the  color  of  the  oxid.    Examine  other  oxids  of  lead. 


78          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

c.    Action  of  Acids  upon  Lead.     Try  the  effect  of   acids, 
hot  and  cold,  upon  lead. 

Hydrogen       Lead 
Lead  +  Nitrogen    =  Nitrogen  +  Hydrogen      N.trogen 


Osygen          Oxygen 

d.  Lead   Chlorid.     To  a  solution  of  lead  nitrate,  add  hy- 
drochloric acid  until  no  further  precipitate  is  formed.     Try 
the  solubility  of  the  resulting  lead  chlorid  in  cold  and  hot 

water. 

Lead  Hydrogen 

__.,  Hydrogen  Lead  „*,    . 

Nitrogen  +     *  =  +  Nitrogen 

Chlorm  Chlorm 
Oxygen  Oxygen 

e.  Lead  Sulfate.     To  a  cold  solution   of  lead  nitrate,  add 
dilute  sulfuric  acid  until  no  further  lead  sulfate  is  formed. 
Try  its  solubility  in  cold  and  hot  water. 

Lead  Hydrogen      Lead  Hydrogen 

Nitrogen  +  Sulfur       =  Sulfur     +  Nitrogen 
Oxygen       Oxygen          Oxygen        Oxygen 

f.  Lead  Sulfid.  Pass  a  stream  of 
hydrogen  sulfid  through  a  solution  of 
lead  nitrate.  State  the  color  of  the 
precipitate. 


Hydrogen     Lead 
Nrtrogen  +  J^     =  Sulfur  +  N.trogen 

Oxygen  Oxygen 

g.  Replacement  of  Lead  by  Zinc.  In 
a  solution  of  lead  nitrate,  place  a  strip 
of  zinc,  and  note  the  formation  of 
metallic  lead  crystals.  The  zinc  evi- 
dently takes  the  place  of  the  lead  in  the  nitrate.  A  very 
pretty  effect  can  be  obtained  by  cutting  the  zinc  so  that  it 
can  be  bent  into  a  tree-like  shape. 

Lead  Zinc 

Nitrogen  +  Zinc   =  Nitrogen  +  Lead 

Oxygen  Oxygen 


ELEMENTS    AND    COMPOUNDS  ?9 

EXPERIMENT   54 
Tin 

a.  Properties.     State  the  properties  of  tin. 

b.  Tin  Oxid.     Oxidize  a  little  tin  in  a  crucible  by  heating, 
stirring  it  with  an  iron  rod. 

Tin  +  Oxygen  = 


c.  Action    of  Hydrochloric   and  Sulfuric   Acids  upon   Tin. 
Try  the  action  of  hydrochloric*  acid  and  sulfuric  acid  on 
tin. 

«)       «.+2£T=  SI*  +**••-• 

Hydrogen      Tin 

(2}  Tin  -f  Sulfur       =  Sulfur    +  Hydrogen 

Oxygen          Oxygen 

d.  Action  of  Nitric  Acid  upon  Tin.     Try  nitric  acid  and  tin. 
In  this  case,  the  nitrate  is  not  formed.     Tin   is  peculiar   in 
that    under   certain    conditions    it  can  form  an  acid.     The 
white  precipitate  you  obtain  is  called  metastannic  acid. 

e.  Replacement  of  Tin  by  Zinc.     In  a  solution  of  tin  chlorid, 
place  a  strip  of  zinc,  and  note  the  metallic  tin  deposited  on 
the  zinc.     Evidently  the  zinc  joins  the  chlorin  left  by  the  tin. 


d.  Tin 


EXPERIMENT    55 

Aluminum 

a.  Properties.     State  the  properties  of  aluminum. 

b.  Does  it  oxidize  readily  ? 

*  A  small  piece  of  copper  with  the  tin  aids  the  action. 


8o        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

c.    Action  of  Adds  upon  Aluminum.     Try  acids,  hydrochlo- 
ric, nitric,  and  sulfuric,  and  verify  the  following  :  - 

Hydrogen      Aluminum 
(i)  Aluminum  +  =  +  Hydrogen 


Hydrogen      Aluminum 

(2)          Aluminum  +  Sulfur       =  Sulfur        +  Hydrogen 
Oyxgen          Oxygen 

d.  Action  of  Aluminum  upon  Alkalis.  In  a  strong  solu- 
tion of  sodium  or  potassium  hydroxid,  place  a  strip  of  alu- 
minum, and  note  the  evolution  of  hydrogen. 

Sodium         Aluminum 

Aluminum  +  Oxygen     =  Oxygen       +  Hydrogen 
Hydrogen      Hydrogen 


EXPERIMENT    56 
Iron 

For  the  properties  and  the  oxid  of  iron,  and  for  the  action 
of   sulfuric  acid  upon    iron,    see   Experiments    14  a,  b,  and 

•s/ 

a.  Action  of  Hydrochloric  Acid  upon  Iron.     Try  the  effect 
of  hydrochloric  acid  upon  iron. 

,    Hydrogen      Iron 
Il0n  +  Chlorin      =  Chlorin  +  H^m^ 

Remark.     Iron    unites  with    two   proportions   of   chlorin, 
forming  salts  called  ferrous  and  ferric  chlorid  respectively. 

b.  Try  iron  and  nitric  acid. 

Hydrogen      Iron 

Iron  +  Nitrogen  =  Nitrogen  +  ^drogen      Nitrogen 
Omen         Oxygen 


ELEMENTS    AND    COMPOUNDS  8 1 

EXPERIMENT    57 
Nickel 

a.  Properties.     State  the  properties  of  nickel. 

b.  Try  to  oxidize  nickel. 

c.  Action  of  Acids  upon  Nickel.     Dissolve  nickel  in  dilute 
nitric  and  sulfuric  acids.     . 

Hydrogen     Nickel 

(,)       Nickel  +  Nitrogen  =  Nitrogen  +  *ydrogen  +  Nltrogen 
Oxygen        Oxygen       Oxygen        Oxygen 

Hydrogen      Nickel 

(2)        Nickel  -I-  Sulfur      =  Sulfur  +  Hydrogen 
Oxygen          Oxygen 

EXPERIMENT   58 
Barium 

There  is  an  element  very  difficult  to  obtain  called  barium. 
It  is  similar  to  calcium  in  its  chemical  properties,  decom- 
posing water  with  the  evolution  of  hydrogen.  Its  com- 
pounds, however,  are  important. 

a.  Barium  Oxid.     Examine  a  little  barium  oxid,  and  note 
its  properties. 

b.  Barium  Hydroxid.     Make  the  hydroxid  from  the  oxid, 
and  test  it  with  its  moist  litmus  paper. 

c.  Barium  Chlorid  and  Nitrate.     Dissolve  a  little  barium 
oxid  in  hydrochloric  acid,  and  crystallize. 

Disolve  a  little  barium  oxid  in  nitric  acid,  and  crystallize. 

Barium        Hydrogen  _  Barium        Hydrogen 
Oxygen       Chlorin     ~  Chlorin       Oxygen 


82        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Hydrogen      Barium 

/  x  Barium  Hydrogen 

(2)  .  +  Nitrogen  =  Nitrogen +_/ 

Oxygen        _  rt  Oxygen 

Oxygen          Oxygen 

d.  Barium  Sulfate.     To  a  solution  of  barium  nitrate,  add 
sulfuric  acid  until  no  further  precipitate  is  formed. 

Barium  Hydrogen  Barium  Hydrogen 
Nitrogen  +  Sulfur  =  Sulfur  +  Nitrogen 
Oxygen  Oxygen  Oxygen  Oxygen 

e.  Flame  Test.     In  a   flame,  hold  a  platinum  wire    upon 
which  there  is  a  little   barium   chlorid,  and   note  the  color. 
This  color  is  characteristic  of  barium  compounds. 


EXPERIMENT    59 

Strontium 

The  element  strontium  is  a  yellow  metal  very  similar  to 
calcium  and  barium  in  its  chemical  properties. 

a.  Strontium  Oxid.     Get  the  properties  of  strontium  oxid. 

b.  Strontium    Hydroxid.       Make    the    hydroxid  from  the 

oxid,  and  test  it  with  litmus. 

Strontium 
Strontium        Hydrogen  = 

Oxygen  Oxygen          Hydn)gen 

c.  Strontium  Chlorid  and  Nitrate.     Make  the  chlorid  and 

nitrate. 

,  v  Strontium      Hydrogen  _  Strontium      Hydrogen 

Oxygen  Chlorin      ~  Chlorin  Oxygen 

Hydrogen      Strontium 
(2)         Stawtam         [      *n  =  mt  +  Hydrogen 

°^en  Oxygen         Oxygen          Oxygen 

d.  Flame  Test.    In  a  flame,  hold   a   platinum   wire   upon 
which  there  is  a  little  strontium  chlorid,  and  note  the  cQlor, 
which  is  characteristic  of  all  strontium  compounds. 


ELEMENTS    AND    COMPOUNDS  83 

EXPERIMENT   60 
Silver 

a.  Properties.     Examine   a  piece   of   silver,  and  state  its 
properties. 

b.  Try  to  oxidize  it  by  heat. 

c.  Action  of  Acids  upon   Silver.     Try  the  effect  of  acids 

upon  silver. 

Hydrogen      Silver 

Silver +  N,trogen  =  N.trogen  +  J      *     +  ^ 

Oxygen          Oxygen 

d.  Silver  Chlorid.     Dissolve  5  grms.  of  silver  nitrate  in  a 
little  water,  and  add  sodium  chlorid  until  no  further  precipi- 
tate is  formed.     Allow  a  part  of  the  precipitated  silver  chlo- 
rid to  stand  in  the  light,  and  try  the  solubility  of  the  rest  in 
ammonium  hydroxid. 

Silver  Sodium 

Nitrogen  +  ^lum  =  **™.     +  Nitrogen 
_  Chlorm       Chlorm        . 

Oxygen  Oxygen 

e.  Silver  Bromid.     Perform    a    similar  experiment,  using 
sodium  bromid  instead  of  sodium  chlorid. 

Silver  Sodium 

Nitrogen  +  fdlum  =  **™.     +  Nitrogen 
„  Bromin        Bromm        . 

Oxygen  Oxygen 

/  Silver  lodid.  Perform  a  similar  experiment,  using  po- 
tassium iodid. 

Silver  Potassium 

_.,  Potassium  Silver       __., 

Nitrogen  +  +  Nitrogen 

»  lodm  lodm         _ 

Oxygen  Oxygen 

g.  Silver  Oxid.  Perform  a  similar  experiment,  using 
sodium  hydroxid.  In  this  case,  silver  oxid  is  formed,  instead 
of  the  hydroxid  that  you  would  naturally  expect. 


84        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Silver          Sodium  Sodium       Hydrogen 

Nitrogen  -f  Oxygen     —  +  Nitrogen  +  Nitrogen 

Oxygen       Hydrogen  ^          Oxygen       Oxygen 

h.  Replacement  of  Silver  by  Copper.  In  a  solution  of  sil- 
ver nitrate,  place  a  strip  of  clean  copper.  Note  the  metallic 
silver  formed. 

Silver  Copper 

Nitrogen  +  Copper  —  Nitrogen  +  Silver 

Oxygen  Oxygen 

/.  Reduction  of  Silver  Chlorid  by  means  of  Hydrogen.  In  a 
hard  glass  tube,  place  a  little  dry  silver  chlorid,  and,  passing 
a  stream  of  hydrogen  over  it,  heat  it  to  redness.  Note  the 
metallic  silver  formed. 

Silver  Hydrogen 

_,.     .    +  Hydrogen  =    *  +  Silver 

Chlonn        J  Chlorm 


EXPERIMENT   61 
Gold 

a.  Properties.     Note  the  properties  of  gold.    Hold  a  piece 
of  gold  leaf  up  to  the  light. 

b.  Action  of  Acids  upon  Gold.       Try  the  effect  of  ordinary 
acids  upon  gold. 

c.  Action  of  Aqua  Regia  upon  Gold.     In  a  test  tube,  place 
3  cc.  of  hydrochloric  acid,  add  i  cc.  of  nitric  acid,  and  warm 
the  mixture.     Note  the  chlorin  evolved.     Place  a  few  pieces 
of  gold  leaf  on  a  watch  glass,  and  add  a  few  drops  of  the 
mixed  acids  (called  aqua  regia) ;  then,  when  all  is  dissolved, 
evaporate  it,  and  obtain  crystals  of  gold  chlorid. 

The  chlorin,  when  in  the  nascent  state  (see  Experiment 
40),  will  unite  with  gold,  whereas,  when  it  is  already  united 
with  hydrogen  in  hydrochloric  acid,  it  has  no  effect. 


ELEMENTS    AND    COMPOUNDS  85 

EXPERIMENT  62 
Platinum 

a.  Properties.     Note  the  properties  of  platinum.     Try  to 
melt  it  before  the  blowpipe. 

b.  Action  of  Adds  upon  Platinum.     Try  the  effect  of  acids 
as  you  did  in  the  case  of  gold. 

c.  Action  of  Metals  upon  Platinum.     On  a  piece  of  char- 
coal, place  a  bit  of  lead  together  with  a  small  piece  of  plati- 
num, and  heat  them  before  the  blowpipe.     Should  metals  be 
heated  in  platinum  crucibles  ? 


PART  II. 
LAWS   AND   THEORIES   OF   CHEMISTRY 


PART    II. 
LAWS   AND    THEORIES   OF   CHEMISTRY 


Introduction.  The  human  mind  has  always  been  more  or 
less  curious  in  regard  to  the  phenomena  of  nature.  In  the 
time  of  the  Greeks,  the  philosophers  tried  to  explain  these 
phenomena  by  speculation  almost  entirely.  Not  until  the 
time  of  Bacon  did  men  become  aware  of  the  fact  that,  in 
order  to  understand  the  changes  that  are  going  on  about  us 
in  the  physical  world,  it  was  necessary  to  make  observations 
of  these  changes,  and  then  arrange  them  in  classes  in  order 
to  see  what  relation  they  bore  to  each  other.  In  this  way 
science  was  born. 

But  it  was  soon  found  that  a  mere  observation  of  facts 
was  not  enough.  The  question  immediately  arose  in  the 
mind  of  each  observer,  "  How  came  it  so  ?  "  In  order  to 
explain  the  facts  observed,  the  necessity  of  some  sort  of  an 
hypothesis  was  seen.  Thus,  when  the  phenomena  of  light 
and  heat  had  been  investigated  by  innumerable  experiments, 
the  questions  arose  :  "  What  is  light  ?  What  is  heat  ?  " 
The  hypothesis  was  made  that  these  phenomena  were  caused 
by  the  vibration  of  an  hypothetical  substance  that  pervaded 
all  space.  Obviously,  this  is  a  mere  guess  ;  but,  so  long  as 
the  facts  agree  with  the  guess,  men  are  justified  in  believing 
that  the  probability  in  favor  of  its  truth  is  great.  Again, 
after  the  facts  of  chemistry  had  been  sufficiently  investigated 
by  experiment,  men  asked  themselves  the  question  :  "  Why 
does  matter  act  so  ? "  The  Atomic  Hypothesis  or  Theory 

89 


QO        AN     ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

was  proposed  to  explain  the  facts.  This  theory  is  so  closely 
in  accord  with  what  has  been  observed  by  experiment,  and 
has  successfully  withstood  so  many  able  assailants,  that  men 
are  agreed  that  its  chance  of  being  the  truth  is  extremely 
great.  However,  should  facts  heretofore  unknown  be  dis- 
covered that  conflict  with  the  atomic  theory,  we  should  be 
forced  to  discard  it  and  look  for  some  other  explanation. 

The  student  must  be  careful  in  his  thinking  to  distinguish 
clearly  between  those  things  which  we  know  from  observa- 
tion and  experiment,  and  those  things  which  we  assume  in 
order  to  explain  what  we  have  observed.  The  statements 
of  facts  learned  by  observation,  we  call  laws.  The  assump- 
tions that  we  make  to  explain  these  facts,  we  call  theories. 

Indestructibility  of  Matter.  The  student  has  probably  by 
this  time  become  convinced,  from  the  experiments  in  Part  I., 
that  matter  is  indestructible.  This  undeniable  fact  may  be 
stated  thus :  Matter  can  neither  be  destroyed  nor  created  by 
any  known  human  agency.  If  this  be  true,  then  it  must  also 
be  true  that  the  sum  of  the  weights  of  the  factors  of  a 
chemical  change  equals  the  sum  of  the  weights  of  the 
products.  Factors  signify  the  substances  put  together, 
and  products  the  substances  obtained. 

EXPERIMENT  i 

The  Sum  of  the  Weights  of  Factors  in  a  Chemical  Change 
equals  the  Sum  of  the  Weights  of  the  Products 

Weigh  out  exactly  5  grms.  of  barium  nitrate  and  3.5 
grms.  of  potassium  sulfate,  and  dissolve  each  portion  in 
100  cc.  of  water  in  separate  beakers.  Heat  the  solutions  to 
boiling,  then  pour  the  potassium  sulfate  into  the  barium 
nitrate.  Weigh  a  filter  paper  carefully  to  i  c.  grin.,  and, 
when  the  precipitate  has  settled,  pour  off  the  liquid  on  to 


LAWS    AND    THEORIES    OF    CHEMISTRY  9 1 

the  filter,  catching  the  filtrate  in  a  carefully  weighed  porce- 
lain dish.  Pour  on  to  the  precipitate,  which  is  still  left  in 
the  beaker,  about  100  cc.  of  hot  water;  then,  after  it  has 
settled,  pour  off  the  liquid  on  to  the  filter  as  before.  Do 
this  twice,  and  then,  by  means  of  a  fine  stream  of  water 


from  a  wash  bottle,  transfer  all  of  the  precipitate  to  the  filter. 
A  camel's  hair  brush  will  aid  greatly  in  getting  the  precipitate 
from  the  sides  of  the  beaker.  If  a  clean  glass  stirring-rod 
is  used  to  guide  the  stream,  the  liquid  will  not  run  down 
the  outside  of  the  beaker.  Then  wash  the  precipitate  on 
the  filter  several  times  with  hot  water  from  the  wash 
bottle.  Dry  the  filter  in  an  oven  at  a  temperature  of 
about  100°,  allow  it  to  cool,  and  weigh  it.  Evaporate 
the  filtrate  to  dryness,  and  weigh.  Arrange  your  calcula- 
tions thus : — 

Wt.  of  filter  paper  -f-  barium  sulfate  = 

Wt.  of  filter  paper 

Wt.  of  barium  sulfate 

Wt.  of  dish  +  potassium  nitrate        = 

Wt.  of  dish 

Wt.  of  potassium  nitrate 

See  whether  the  experiment  shows  the  truth  of  the  law. 


92          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Law  of  Definite  Proportions  by  Weight.  The  next  ques- 
tion that  naturally  arises  is,  "  Does  chemical  action  take 
place  between  definite  quantities  of  elements,  or  are  these 
quantities  variable  ?  "  It  has  been  found  by  innumerable 
experiments  that :  —  Every  chemical  compound  has  a  fixed  and 
unalterable  composition.  This  is  known  as  the  Law  of  Definite 
Proportions  by  Weight. 


EXPERIMENT    2 


Law  of  Definite  Proportions  by  Weight 

Weigh  a  porcelain  boat  accurately  on  a  delicate   balance. 
Place  in  it  about  i  grm.  of  c.p.  cupric  oxid  and  weigh  it 


again.  Then  place  the  boat  with  contents  in  a  glass  tube 
containing  stoppers,  with  entrance  and  exit  tubes.  Bend  a 
piece  of  copper  foil  so  that  it  will  fit  halfway  around  the 
tube  for  its  entire  length.  Connect  the  tube  with  a  hydrogen 
generator,  and  allow  a  stream  of  hydrogen,  dried  by  being 
passed  through  sulfuric  acid,  to  pass  through  the  apparatus. 
When  all  the  air  has  been  driven  out,  heat  the  tube  with  a 
Bunsen  flame  until  all  the  cupric  oxid  is  reduced  to  copper. 
Allow  the  tube  and  contents  to  cool  in  the  stream  of  hydro- 


LAWS    AND    THEORIES    OF    CHEMISTRY.  93 

gen,  and  weigh  it  again.     In  another  weighed  porcelain  boat, 
place  about  2  grms.  of  the  same  oxid,  and  weigh  again  care- 
fully.     Go    through    the    same    operation,    and   obtain    the 
weight  of  the  boat  and  copper  when  all  is  cool. 
Arrange  your  calculations  as  follows  : 

(a)    Wt.  of  boat  and  oxid  (b)  Wt.  of  boat  and  oxid      = 

Wt.  of  boat  Wt.  of  boat 

Wt.  of  copper  obtained  =  Wt.  of  copper  obtained  = 

Then  make  the  following  proportion  :  — 

Wt.  of  oxid  in  a     :      Wt.  of  oxid  in  b     : :     Wt.  of  copper 
obtained  in  a       :      Wt.  of  copper  obtained  in  b. 

Law  of  Multiple  Proportions  by  Weight.  It  will  be  seen 
that  there  are  some  elements  that  unite  to  form  more  than 
one  compound,  as  for  instance  sulfur  and  oxygen,  carbon  and 
oxygen,  and  carbon  and  hydrogen.  Analysis  shows  that 
these  compounds  verify  the  law  of  definite  proportions. 
The  question  arises :  "  Is  there  any  fixed  relation  between 
the  comparative  amounts  of  the  same  element  in  such  com- 
pounds ?  "  By  experiment,  chemists  have  proved  that : 

When  two  elements  combine  to  form  more  than  one  compound, 
the  amounts  of  one  of  the  elements  which  combine  with  a  fixed 
amount  of  the  other  bear  to  each  other  a  simple  ratio. 

This  is  called  the  Law  of  Multiple  Proportions. 


EXPERIMENT    3 
Law  of  Multiple  Proportions 

There  are  four  compounds  of  copper  and  oxygen,  and,  by 
experiments  similar  to  the  preceding,  chemists  have  found 
that  the  compounds  have  the  following  composition  :  — 


94        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Copper  Oxygen 

1.  94-09%  5-9i% 

2.  '  88.83%  ".17% 

3.  79-90%  20.10% 

4-  66.52%  33-47% 

Find  in  each  of  the  last  three  cases  what  amount  of 
oxygen,  according  to  the  law  of  definite  proportions,  would 
unite  with  94.09  parts  of  copper.  Calculate  thus  in  each 

94.09     :    88.83     :  :    x    :     11.17. 

Find  out  whether  the  four  amounts  of  copper  in  the  differ- 
ent cases  bear  to  each  other  a  simple  ratio. 

Quantitative  Analysis.  In  Experiment  2,  the  student 
has  performed  two  quantitative  analyses.  It  is  by  various 
methods  that  the  chemist  is  able  to  determine  the  percentage 
composition  of  compounds.  It  will  perhaps  be  well  for  the 
student  to  analyze  a  .compound  by  precipitating  one  of  the 
elements  in  combination  with  some  other  element,  the  per- 
centage composition  of  the  precipitated  compound  being 
first  determined. 

EXPERIMENT    4 
Quantitative  Analysis  of  Sodium  Chlorid 

a.  In  order  to  analyze  sodium  chlorid,  it  is  necessary  to 
know  the  percentage  composition  of  silver  chlorid.  To  find 
this,  proceed  as  follows :  — 

Weigh  out  on  a  delicate  balance  about  .5  grm.  of  pure  silver, 
recording  the  exact  weight.  Dissolve  the  silver  in  a  little 
c.p.  nitric  acid  diluted  with  water  in  a  clean  porcelain  dish, 
and  evaporate  it  to  dryness.  Dissolve  the  crystals  in  about 
200  cc.  of  distilled  water.  You  now  have  a  solution  of  silver 
nitrate,  which  contains  all  the  silver.  Make  a  solution  of 
about  .5  grms.  of  c.p.  sodium  chlorid  in  100  cc.  of  dis- 


LAWS    AND    THEORIES    OF    CHEMISTRY  95 

tilled  water.  Add  from  this  to  the  silver  nitrate  solution  as 
long  as  any  precipitate  is  formed,  and  boil.  When  the  pre- 
cipitate has  settled,  add  a  drop  or  two  of  the  sodium  chlorid, 
to  make  sure  that  all  the  silver  is  precipitated.  Filter 
off  the  clear  liquid  while  it  is  hot.  Add  about  300  cc.  of  hot 
distilled  water  to  the  remaining  precipitate,  stir,  allow  it  to 
settle,  and  filter  it  again.  Do  this  twice,  then  throw  all  the 
precipitate  on  the  filter,  being  sure  to  lose  none.  Wash  it  on 
the  filter  two  or  three  times  with  hot  water  from  the  wash 
bottle,  and  allow  it  to  dry  in  an  oven  at  a  temperature  of  about 
100°.  When  the  precipitate  is  dry,  ^remove  as  much  of  it  as 
possible  to  a  weighed  porcelain  crucible.  Burn  the  paper, 
holding  it  with  a  platinum  wire  over  a  clean  porcelain 
plate.  Do  not  let  the  wire  come  in  contact  with  any  of  the 
precipitate.  When  all  the  carbon  from  the  paper  is  burned 
away,  transfer  all  the  residue  by  means  of  a  brush  to  the 
crucible  cover.  Some  of  the  silver  chlorid  on  the  filter  will 
have  been  reduced  to  silver  by  the  burning  paper.  To 
change  this  back  to  silver  chlorid,  add  two  or  three  drops  of 
nitric  acid,  warm,  then  add  two  of  three  drops  of  hydrochlo- 
ric acid,  and  evaporate  it  to  dryness.  Now  barely  melt  the 
chlorid  in  the  crucible,  allow  it  to  cool  in  a  desiccator,  and 
weigh  the  whole. 

Arrange  your  calculations  as  follows  :  — 

Wt.  of  silver  chlorid  and  crucible  =  m 

Wt.  of  silver  taken  =  a 

Wt.  of  crucible  =  n 

Wt.  of  silver  chlorid  =  m  —  n    = 

Percentage  of  silver  in  silver  chlorid  = = . 

m  —  n 

Then  i  oo  %  -  -  %  =  percentage  of  chlorin  in  silver 

chlorid. 

Your  final  result  should  be  somewhere  near  24.7%. 


96        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

b.  Now  heat  about  10  grms.  of  c.p.  sodium  chlorid  in  a 
porcelain  crucible,  to  drive  off  any  absorbed  moisture.  Cool 
it  in  a  desiccator.  Then  weigh  out  on  the  delicate  balance 
about  .5  grin,  of  this,  and  dissolve  it  in  about  100  cc.  of  dis- 
tilled water  in  a  clean  beaker.  Heat,  and  add  from  a  silver 
nitrate  solution  until  all  the  chlorin  is  precipitated  as  silver 
chlorid.  Proceed  exactly  as  in  the  preliminary  part  of  this 
experiment,  and  obtain  the  weight  of  the  silver  chlorid. 

Arrange  your  calculations  thus  :  — 

Wt.  of  sodium  chlorid  taken  =  a 

Wt.  of  crucible  and  silver  chlorid  =  b 

Wt.  of  crucible  =  c 

Wt.  of  silver  chlorid  =  b  —  c  = 

Since  silver  chlorid  contains  24.73%  °f  chlorin,  there  are 
24.73  (f>  —  c)  grms.  of  chlorin  in  the  silver  chlorid  obtained. 
But  this  chlorin  came  from  the  a  grms.  of  sodium  chlorid 
taken  ;  therefore  the  percentage  of  chlorin  in  sodium  chlorid 

is  ^73_I       -I,     Your  result  should  be  about  60.6%. 

Atomic  Theory.  As  soon  as  the  preceding  facts,  i.e.,  the 
indestructibility  of  matter,  the  law  of  definite  proportions, 
and  the  law  of  multiple  proportions,  had  been  established, 
men  of  science  immediately  began  to  look  for  some  explana- 
tion of  them.  It  remained  for  Dalton  in  the  early  part  of 
this  century  to  propose  an  hypothesis  that  is  accepted  to  the 
present  day.  We  are  compelled  to  believe  in  its  truth,  for 
we  can  think  of  no  other  hypothesis  that  will  account  for  the 
facts. 

Dalton 's  hypothesis  is  this.  All  matter  is  composed  of 
minute  particles,  which  cannot  be  divided  by  any  chemi- 
cal means.  Each  of  these  particles  of  any  simple 
substance  (i.e.  element)  is  like  every  other  one  of  that  sub- 


LAWS    AND    THEORIES    OF    CHEMISTRY  Q/ 

stance,  both  in  properties  and  weight.  These  particles  are 
called  atoms.  Therefore,  an  atom  is  the  smallest  particle 
of  matter  that  can  enter  into  chemical  combination.  These 
like  or  unlike  atoms  may  unite  with  each  other,  and  form 
other  particles  which  are  called  molecules.  Therefore,  a 
molecule  is  the  smallest  particle  of  matter  that  can  exist  and  still 
have  all  the  properties  of  the  substance,  and  is  made  up  of  atoms 
chemically  united*  This,  briefly  stated,  is  Dalton's  Atomic 
Theory.  It  must  be  borne  in  mind  that  it  is  simply  an 
hypothesis  or  guess  formulated  to  explain  known  facts.  If 
the  hypothesis  is  the  actual  truth,  then  it  follows  that,  in  the 
case  of  copper  and  oxygen,  when  their  atoms  are  chemically 
united  to  form  copper  oxid,  the  proportions  will  always  be 
the  same.  In  the  case  of  two  compounds  of  the  same  ele- 
ments, as  for  instance  the  two  oxids  of  copper,  it  follows 
that  the  amount  by  weight  of  oxygen  in  the  first  must  bear 
a  simple  ratio  to  the  amount  by  weight  of  oxygen  in  the 
second.  For  evidently  twice  as  many  atoms  of  oxygen 
would  unite  with  a  definite  number  of  atoms  of  copper  to 
form  the  second  oxid  as  would  unite  with  the  same  number 
of  atoms  of  copper  to  form  the  first  oxid.  In  all  cases,  the 
ratio  must  be  a  simple  one,  since,  if  matter  is  made  up  of 
atoms,  chemical  combination  must  take  place  between  the 
whole  atoms. 

Atmospheric  Pressure.  Barometer.  In  order  to  understand 
the  full  meaning  of  the  atomic  theory,  it  will  be  necessary 
for  the  student  to  perform  a  number  of  experiments  with 
gases.  To  do  this,  he  must  understand  the  meaning  of 
atmospheric  pressure,  and  the  effects  of  changes  of  pressure 
and  of  temperature  on  the  volumes  of  gases.  First  let  us 
take  up  the  subject  of  atmospheric  pressure. 

*  Some  molecules  are  made  up  of  single  atoms. 


98          AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

EXPERIMENT    5 
Barometer 


Take  a  glass  tube  of  uniform 
bore  (5  mm.)  about  a  meter 
long,  and  close  one  end.  Fill  it 
with  mercury,  and,  by  means 
of  a  wire,  remove  what  air  bub- 
bles cling  to  the  sides  of  the  tube. 
Invert  the  tube  in  a  trough  of 
mercury,  and  measure  the  height 
of  the  column.  Such  an  in- 
strument as  this  is  called  a  ba- 
rometer. Now  since  the  mercury 
remains  in  the  tube,  some  force 
must  be  exerted  to  hold  it  there, 
and  the  only  thing  that  can  do 
so  is  the  air  pressing  down 
upon  the  surface  of  the  mercury 
in  the  trough.  Therefore  the 
weight  of  the  mercury  in  the 
tube  balances  a  column  of  air  of 
equal  cross  section  extending  to 
the  top  of  the  atmosphere.  If 
you  should  perform  this  experi- 
ment on  different  days,  you 
would  observe  that  the  column 
stands  at  various  heights. 
Hence  we  see  that  the  pressure 
or  weight  of  the  atmosphere 
resting  upon  a  certain  part  of 
the  earth's  surface  varies  at 


LAWS    AND   THEORIES    OF    CHEMISTRY  99 

different  times.  The  standard  pressure  is  the  pressure  that 
will  hold  up  a  column  of  mercury  760  mm.  high,  when  the 
thermometer  stands  at  o°  C.  This  is  about  15  Ibs.  to  the 
square  inch  or  1033.6  grms.  to  the  square  centimeter. 

Boyle's  Law.  Now  we  are  ready  to  investigate  the  effect 
of  pressure  upon  the  volume  of  a  gas.  We  all  know  by 
every  day  experience  that  gases  can  be  reduced  to  a  smaller 
volume  by  increasing  the  pressure  upon  them,  and  that  they 
expand  when  the  pressure  is  removed  ;  but  most  of  us  are 
perhaps  ignorant,  unless  we  have  studied  physics,  of  the 
exact  effect  of  the  pressure  upon  the  volume.  The  follow- 
ing experiment  will  show  that  the  volume  of  a  gas  varies 
inversely  as  the  pressure  exerted  upon  it ;  i.  e.,  if  you  double 
the  pressure,  you  divide  the  volume  by  two,  or  if  you  treble 
the  pressure,  you  divide  the  volume  by  three. 

Expressing  the  law  algebraically,  we  have 

V  :  V  :  :  P'  :  P 

in  which  V  is  the  volume  at  the  pressure  P,  and  V  is  the 
volume  at  the  pressure  P'.  This  is  called  Boyle  V  Law. 


EXPERIMENT    6 

Law  of  Boyle 

Have  ready  a  clean  dry  glass  tube  of  about  6  mm.  bore, 
closed  at  one  end  and  bent  so  as  to  form  a  narrow  letter  J, 
the  hook  being  about  30  cm.  long.  Let  the  long  arm  be 
about  100  cm.  in  length.  Into  this,  pour  a  little  clean  mer- 
cury, so  that  it  stands  about  the  same  height  in  both  arms. 
It  will  do  no  harm  if  it  stands  a  little  higher  in  the  long  arm. 
Now  fasten  the  tube  to  a  perpendicular  support,  and  meas- 
ure from  the  base  of  the  support  to  the  top  of  the  bore  in 


IOO       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


the  short  arm,  reading 
the  scale  to  tenths  of 
millimeters.  The  edge  of 
a  visiting  card  held  across 
the  tube  will  aid  the  eye 
in  taking  the  reading. 
Then  measure  the  height 
of  the  mercury  in  the  long 
and  short  arms.  Do  not 
touch  the  column  of  con- 
fined air  with  the  hand  ; 
for,  in  that  case,  the 
warmth  of  the  hand  will 
cause  the  air  to  expand 
and  thus  give  incorrect 
readings  on  the  scale. 
Now  take  the  reading 
of  the  barometer.  Four 
into  the  tube  enough  mer- 
cury to  make  the  column 
in  the  long  arm  stand 
about  10  cm.  higher.  .To 
remove  any  air  bubbles 
that  may  be  held  im- 
prisoned by  the  mercury, 
push  a  long  iron  wire  down 
into  the  mercury  ;  then 
withdraw  it,  at  the  same 
time  turning  it  around  in 
the  tube.  Take  the  meas- 
urements as  before.  Do 
this  as  many  times  as  the 
length  of  the  tube  permits. 


LAWS    AND   THEORIES  '  OF"CHEM/STk%  '        : /«TQI 
Arrange  your  numbers  in  a  table  thus  :  — 


£^S             i£ 

^  £               ^              S  1  3       1       £  ai 

^  §*E  •    ;     i~  *   • 
|c|<        ri- 

"Sj^ 

0      . 

jl 

yit 

~~  tc-  ^ 

^=  c  c^1 

^+ 

i== 

ll!      !     II 

rt 

1 

til" 

=  52 

I1 

Hll 

|!« 

1 

Compare  the  volume  at  any  pressure  with   the  volume  at 
any  other  pressure,  and  verify  the  proportion  V  :  V  :  :  P'  :  P. 


EXPERIMENT    7 
Law  of  Charles 

Have  ready  a  tube  of  i  mm.  bore  about  100  cm.  long, 
bent  at  right  angles  at  a  distance  of  about  50  cm.  from  the 
closed  end,  and  containing  a  column  of  dry  air  confined  by 
means  of  mercury,  the  mercury  extending  somewhat  beyond 
the  bend.  Insert  this  through  a  cork  stopper  in  a  hole  in 
the  end  of  a  shallow  tin  tray  containing  melting  ice.  Allow 
the  air  column  to  remain  in  the  ice  for  a  few  minutes, 
keeping  the  ends  of  the  mercury  column  at  the  same  level. 


'  AN'  '  EI5EMENTAHY    EXPERIMENTAL    CHEMISTRY 


Mark  the  end  of  the  air  column,  and  measure  with  a  meter 
rod.  Now  insert  the  confined  column  of  air  in  a  steam 
jacket  containing  a  thermometer  ;  then  allow  steam  to  pass 
through  the  apparatus.  Keeping  the  two  ends  of  the  mer- 
cury level,  again  mark  the  end  of  the  air  column.  Remove  it 


from  the  jacket,  and  measure  it.  The  difference  between  the 
lengths  in  ice  and  in  steam  will  be  the  amount  by  which  the 
air  column  has  expanded  in  being  heated  100°. 

To  find  the  amount  one  centimeter  expands  for  one  de- 
gree, divide  this  result  by  the  original  length  x  the  number 


of  degrees  heated.  This  should  give  a  number  near  .00366, 
which  reduced  to  a  fraction  equals  about  ^\^.  We  there- 
fore see  that,  for  every  degree  it  is  heated  at  a  constant  pres- 
sure, a  volume  of  air  expands  -%\^  of  what  its  volume  was 
at  o°  C.  This  is  known  as  the  Law  of  Charles. 


LAWS    AND    THEORIES    OF    CHEMISTRY  IO3 

Theoretically  then,  the  volume  of  a  gas  at  273°  below 
zero  would  become  nothing.  Of  course  this  is  impossible, 
but  the  point  —  273  is  taken  as  the  absolute  zero  of  temper- 
ature. Then  a  °  in  the  ordinary  scale  wrould  be  a  -f-  273° 
in  the  absolute  scale. 

Deduction  of  the  Formula. 
Let  A  be  the  volume  of  a  gas  at  o°, 

then  A  H  ---  will  be  the  volume  at  i° 

273 

A  +  —  will  be  the  volume  at  2° 
273 

A  +  3—  will  be  the  volume  at  3° 
273 

A  +  —  will  be  the  volume  at  t° 
273 

and  A  +  -  -  will  be  the  volume  at  t'° 

273 

Representing  the  volume  at  t°  by  V,  and  the  volume  at 
t'°  by  V,  we  have 

V  =  A  +  —  ,  andV'  =  A  +  — 
273  '  273 

Dividing  and  canceling  the  A's,  we  have 
V 


v~  273  +r 

Combining  this  with  the  formula  for  pressure,  V  :  V'  :  : 

P'  :  P,  we  obtain 

VP          V'P' 

273  -ft  7  273  +t' 

Boyle's  and  Charles  s  Laws  are  General.  In  the  experi- 
ments we  have  used  air,  but  it  will  be  found  by  trial  that  other 
gases  act  in  the  same  way  as  air.  Thus  we  see  that  differ- 
ent kinds  of  matter  in  the  gaseous  state  have  a  property  in 
common. 


IO4       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Examples,  i.  The  volume  of  a  gas  at  765  mm.  pressure 
and  20°  C.  is  450  cc. ;  what  will  its  volume  be  at  o°  and 
7 60  mm.  ?  Ans.  422.  -|- 

2.  At  what  temperature,    pressure  remaining  the  same, 
will  a  gas  be  double  its  volume  at  o°  ?     Ans.  273°. 

3.  Reduce  the  following  volumes  with  the  annexed  tem- 
peratures and  pressures  to  volumes  at  second  temperatures 
and  pressures :  — 

V  t  P  V      t'  P' 

500       100        800  mm.  ?       20.       650  mm. 

1000        70        900  mm.  ?        o.       200  mm. 

EXPERIMENT  8 
Weight  of  a  Liter  of  Air 

Fit  a  strong  glass  bottle  of  about  one-liter  capacity,  with  a 
one-holed  rubber  stopper  containing  a  glass  tube.  Wire  on 
to  the  tube  a  short  piece  of  antimony-rubber  tubing  fitted 
with  a  pinch  cock.*  Make  all  connections  air-tight  with  vase- 
line. Weigh  the  bottle  and  connections  carefully  to  one 
centigram.  Attach  an  air  pump  to  the  rubber  tube,  pump 
out  as  much  air  as  possible,  close  the  tube  with  the  pinch 
cock,  and  weigh  the  bottle  carefully  again.  Open  the  tube 
under  water  of  about  the  same  temperature  as  the  room. 
After  the  water  has  ceased  rushing  in,  hold  the  bottle  so 
that  the  water  stands  at  the  same  level  both  inside  and  out- 
side, close  the  pinch  cock,  remove  the  bottle,  and  weigh 
it  again.  Take  the  readings  of  thermometer  and  barometer. 

Arrange  your  calculations  thus : 

Wt.  of  bottle  full  of  air  =  a 

Wt.  of  bottle,  air  pumped  out  =   b 

Wt.  of  air  pumped  out  =   a  —  b 

Wt.  of  bottle  with  water  =   d 

Wt.  of  water  =  vol.  air  pumped  out  =  d  —  a  =  e. 


LAWS    AND    THEORIES    OF    CHEMISTRY 


105 


Therefore  e  cc.  of  air  at  the  temperature  and  pressure 
on  the  day  of  the  experiment  weigh  a  —  b  grams.  Calcu- 
late the  weight  of  one  liter.  You  should  obtain  with  this 
apparatus  somewhere  near  1.2  or  1.3  grams. 

The  correct  weight  of  a  liter  of  air  at  o°  and  760  mm. 
pressure  is  1.293  grams,  which  number  we  shall  hereafter  use. 


EXPERIMENT    9 
Weight  of  a  Liter  of  Carbon  Dioxid  at  o°  and  760  mm. 

Fit  a  dry  one-liter  flask  with  a  two-holed  rubber  stopper. 
In  the  holes,  insert  glass  tubes  bent  at  right  angles,  one 
reaching  to  the  bottom  of  the  flask  and  the  other  just 


through  the  stopper.  Fit  with  rubber  tubes  and  pinch  cocks. 
Make  all  joints  air-tight  with  vaseline,  and  weigh  carefully  to 
one  centigram.  Allow  carbon  dioxid,  dried  by  being  passed 
through  sulfuric  acid,  to  pass  through  the  flask  until  it  is  com- 
pletely rilled.  Close  the  tubes  and  disconnect  them.  Open 


IO6       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

one  pinch  cock  for  a  moment  to  relieve  the  pressure,  close  it 
again,  and  weigh  the  flask.  Take  the  temperature  near  the 
flask  and  the  reading  of  the  barometer.  The  increase  in 
weight  will  be  the  amount  by  which  the  weight  of  the  flask 
full  of  carbon  dioxid  exceeds  the  weight  of  'the  flask  full  of 
air.  Mark  the  capacity  of  the  flask,  fill  it  with  water  to  the 
mark,  and  weigh.  Since  i  cc.  of  water, weighs  one  gram,  the 
number  of  grams  of  water  in  the  flask  equals  the  number  of 
cubic  centimeters  in  the  volume  of  the  flask.  Calculate 
what  this  volume  of  air  at  the  temperature  and  pressure  at 
the  time  of  the  experiment  would  be  at  o°  and  760  mm. 
Knowing  that  one  liter  of  air  at  o°  and  760  mm.  weighs 
1.293  grams,  find  the  weight  of  the  air.  Add  to  this  the  in- 
crease in  weight  due  to  the  carbon  dioxid,  and  obtain  the 
weight  of  the  same  number  of  cc.  of  the  gas.  From  this 
result,  obtain  the  weight  of  1000  cc.,  i.  e.,  i  liter  of  carbon 
dioxid.  You  should  obtain  a  number  somewhere  near  1.97. 

EXPERIMENT   10 
Weight  of  a  Liter  of  Hydrogen 

Perform  a  similar  experiment  with  hydrogen  gas.  Use 
a  100  cc.  flask,  and  weigh  it  on  a  delicate  balance.  Keep  the 
flask  inverted  when  filling,  and  be  sure  the  hydrogen  is  dry. 
To  insure  this,  pass  the  hydrogen  not  only  through  sulfuric 
acid  but  also  through  a  U  tube  containing  pieces  of  granu- 
lar calcium  chlorid.  Make  the  hydrogen  with  zinc  and  sul- 
furic acid  (one  part  acid  to  five  of  water).  In  this  case,  the 
weight  of  the  flask  full  of  hydrogen  is  less  than  that  of  the 
flask  full  of  air  ;  hence  you  must  subtract  the  decrease  from 
the  calculated  weight  of  the  flask  full  of  air  at  o°  and  760 
mm.  You  should  obtain  a  result  somewhere  near  .09  grams 
for  i  liter.  (See  illustration  on  page  107.) 


LAWS    AND    THEORIES    OF    CHEMISTRY  IO/ 

Density.  By  the  density  of  an  element  or  compound  is 
meant  its  weight  in  gaseous  form  compared  with  the  weight 
of  an  equal  volume  of  hydrogen,  the  temperature  and  pres- 
sure being  the  same. 


EXPERIMENT    n 

Calculate  the  density  of  air  and  carbon  dioxid  compared 
with  hydrogen,  and  show  that  they  are  14.43  an(^  22  respec- 
tively. 

EXPERIMENT    12 
Density  of  a  Liquid  in  the  form  of  Vapor 

Have  ready  a  clean  dry  100  cc.  flask  fitted  with  a  one- 
holed  rubber  stopper  containing  a  glass  tube  drawn  out  to  a 
diameter  of  about  i  mm.  Make  it  air-tight  with  a  little  vase- 
line, and  weigh  it  carefully  on  a  delicate  balance.  Take  the 
reading  of  the  thermometer  and  barometer.  Place  about 
20  cc.  of  alcohol  in  the  flask,  and  put  the  flask  into  boiling 
water  up  to  its  neck.  When  the  alcohol  has  boiled  away,* 
*  Do  this  gradually  so  as  not  to  blow  out  the  stopper. 


IO8       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


ignite  the  vapor  issuing  from  the  tube,  and  the  moment  that 
the  flame  dies  away,  seal  the  tube  with  a  blowpipe.  Now 
weigh  the  flask  again.  Fill  the  flask  with  water  by  breaking 
the  end  of  the  tube  under  water.  If  the  flask  does  not  fill 
almost  completely,  the  experiment  must  be  repeated.  The 
weight  of  the  flask  full  of  water  minus  the  weight  of  the  flask 
gives  the  volume. 

The  increase  in  weight  will  be  the  weight  by  which  the 
flask  full  of  vapor  of  alcohol  at  100°  exceeds  the  weight  of 
the  flask  full  of  air  at  the  temper- 
ature of  the  room.  Find  what  the 
volume  of  the  air  in  the  flask  would 
become  at  o°  and  760  mm.  Knowing 
the  weight  of  a  centimeter  of  air,  find 
the  weight  of  this  volume  of  air. 
Add  the  increase  in  weight  due  to 
the  alcohol  vapor,  and  obtain  the 
weight  of  the  alcohol  vapor  at  100° 
and  the  pressure  at  the  time  of  the 
experiment.  Calculate  the  volume  of 
the  air  in  the  flask  at  100°.  Let  us 
call  this  M.  Then  if  V  is  the  volume 
of  the  flask  at  the  temperature  of  the  experiment,  the  weight 

of  the  air  left   in   the  flask   at    100°  would  be     -  times  its 

M 

weight  at  the  temperature  of  the  experiment.  We  now  have  the 
weights  of  the  flask  full  of  air  and  alcohol  at  100°.  Hence 
the  density  of  alcohol  referred  to  air  is  easily  found.  But  air 
is  14.43  times  as  heavy  as  hydrogen;  therefore,  to  find  the 
density  of  alcohol  compared  with  hydrogen,  multiply  by  14.43. 

Weights  of  the  Atoms.    The  student  will  remember  that,  in 
the  atomic  theory,  it  is  assumed  that  the  atoms  of  each  ele- 


LAWS    AND    THEORIES    OF    CHEMISTRY  1 09 

ment  are  of  the  same  weight.  Under  the  present  condition 
of  methods  of  investigation,  it  is  impossible  to  determine  the 
absolute  weight  of  these  atoms  with  any  satisfactory  degree 
of  accuracy.  However,  as  soon  as  the  atomic  theory  was 
accepted,  men  began  to  try  to  find  out  the  weight  of  these 
atoms  in  terms  of  some  atom  taken  as  a  standard.  Hydro- 
gen being  the  lightest  element  known,  its  atom  was  taken  as 
this  standard.  It  required  a  number  of  years  and  the  estab- 
lishment of  auxiliary  principles  to  fix  even  these  numbers 
with  any  degree  of  certainty.  As  time  went  on,  investigation 
brought  out  the  principles  of  chemical,  electrical,  thermal,  and 
isomorphic,  equivalence  of  elements.  The  Law  of  Definite 
Properties  by  Volume,  Avogadro  's  Law,  and  the  Periodicity  of 
the  Elements,  were  also  established  :  and  with  the  aid  of  these 
it  has  been  possible  to  determine  the  relative  weights  of  the 
atoms  to  a  remarkable  degree  of  accuracy. 

We  will  take  up  these  subjects  in  turn. 

Chemical  Equivalence.  By  experiment,  it  is  found  that  on 
causing  hydrogen  and  chlorin  to  unite,  the  ratio  of  the  parts 
by  weight  are  i  :  35.4.  That  is,  the  chemical  value  of 
chlorin  compared  with  hydrogen  is  35.4.  In  like  manner,  8 
parts  by  weight  of  oxygen  unite  with  i  part  by  weight  of 
hydrogen.  We  therefore  say  the  chemical  equivalence  of 
oxygen  is  8.  This  is  also  called  its  combining  number.  To 
find  the  chemical  equivalent  of  an  element  in  terms  of 
another,  it  is  necessary  only  to  find  the  amount  by  weight 
of  the  first  that  unites  with  a  fixed  amount  of  the  second,  or 
the  amount  which  replaces  a  fixed  quantity  in  a  compound. 
In  some  cases,  that  is,  where  there  is  more  than  one  com- 
pound between  the  same  element,  it  is  difficult  to  decide 
which  number  to  take.  If  we  knew  the  exact  number  of 
atoms  of  each  kind  that  unite  to  form  a  compound,  or  the 
exact  number  of  atoms  of  an  element  replaced  in  a  compound 


IIO        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

by  one  atom  of  another  element,  it  would  be  easy  to  deter- 
mine the  atomic  weight  directly  from  the  chemical  equiva- 
lence. Thus,  if  we  knew  that  one  atom  of  hydrogen  united 
with  one  atom  of  oxygen  to  form  water,  we  could  be  sure 
that  the  atomic  weight  of  oxygen  was  8.  Or,  if  we  knew 
that  two  atoms  of  hydrogen  united  with  one  of  oxygen,  we 
could  be  sure  of  the  number  16  as  the  atomic  weight  of 
oxygen.  Not  knowing  this,  however,  we  see  that  our  data 
is  insufficient,  and  we  must  look  farther. 


EXPERIMENT    13 
Chemical  Equivalence  of  Zinc 

Have  ready  a  clean  100  cc.  flask  containing  a  dilute  solution 
of  c.p.  hydrochloric  acid  (10  cc.  of  water  to  20  cc.  of  acid). 
Fit  the  flask  with  an  air-tight  one-holed  rubber  stopper  fitted 
with  a  delivery  tube.  Weigh  out  as  nearly  as  possible  1.25 

grms.  of  c.p.  zinc,  recording 
the  exact  weight.  Place  in  a 
pneumatic  trough. a  500  cc. 
flask  inverted  and  full  of 
water.  Have  the  delivery 
tube  clamped  so  that  its  end 
is  under  the  mouth  of  the 
flask.  Remove  the  stopper 
from  the  100  cc.  flask,  drop 
the  zinc  into  it,  close  it 

quickly,  and  catch  all  the  evolved  hydrogen.  When  all  the 
zinc  is  dissolved,  hold  the  flask  containing  the  hydrogen  so 
that  the  water  stands  at  the  same  level  both  inside  and  out- 
side. Do  not  touch  with  the  hand  that  part  of  the  flask  con- 
taining the  hydrogen.  Take  the  temperature  near  the  flask 


LAWS    AND    THEORIES    OF    CHEMISTRY 


II  I 


and  the  reading  of  the  barometer.  Mark  the  point  on  the 
flask  at  which  the  gas  stands,  remove  it,  and  fill  it  with  water 
to  this  point.  Then  weigh  the  flask  and  water.  From  this, 
obtain  the  volume  of  the  gas. 

To  obtain  the  best  results,  it  will  be  necessary  to  take  into 
account  the  fact  that,  besides  the  hydrogen,  there  is  vapor  of 
water  in  the  flask,  which  helps  bear  the  pressure  of  the  air 
upon  the  surface  of  the  water  in  the  trough.  This  varies  for 
different  temperatures.  The  following  table  is  near  enough 
for  all  practical  purposes. 


DEGKKUS. 

X 

H   X 
H  x  p 

Gfl 

1      jj^l 

1 

Hi 

a 
a   > 

j  *jj 

i 

X 

MlLI.IMICTKKS 
OF 
MliKCDKY. 

IO 

9-2 

16     13.6 

22 

19.7 

28 

28.1 

II 

9.8 

17     14.4 

=3 

20.9 

29 

29.8 

.I2 

10.5 

18     154 

24 

22.2 

3° 

31-5 

13 

II, 

19     16.4 

=5 

23-5 

31 

33-4 

14 

II.9 

20       17.4 

26 

25.0 

32 

35-4 

15 

12.7 

21       18.4 

27 

26.5 

33 

37-4 

Subtract  the  number  of  millimeters  of  mercury  in  the  table 
at  the  temperature  corresponding  to  the  temperature  observed 
near  the  flask,  and  you  will  have  the  true  pressure  to  which 
the  hydrogen  is  subjected. 


112       AN    ELEMENTARY    EXPERIMENTAL     CHEMISTRY 

Calculate  by  means  of  the  formula 
V  P  V  P' 


273  +  t         273  +  t' 

what  the  volume  of  the  hydrogen  would  be  at  o°  and  760 
mm.  Knowing  that  i  cc.  of  hydrogen,  at  o°  and  760  mm., 
weighs  .00009  grins.,  calculate  the  weight  of  the  hydrogen ; 
then  obtain  from  this  the  number  of  grams  of  zinc  that  will 
replace  one  gram  of  hydrogen.  You  should  obtain  some- 
where near  the  number  32.5. 

Should  time  allow,  it  will  be  interesting  for  the  student  to 
determine  the  number  for  iron,  using  piano  wire  instead  of 
zinc.  The  number  in  this  case  is  27.9. 

Electrical  Equivalents.  Michael  Faraday  found  that,  when 
a  current  of  electricity  was  passed  through  a  substance  which 
it  decomposed,  the  quantity  of  the  substance  decomposed 
varied  with  the  strength  of  the  current.  If  the  same  current 
passes  through  two  substances,  such  as  a  solution  of  copper 
sulfate  and  a  solution  of  silver  nitrate,  the  amount  of  copper 
deposited  is  to  the  amount  of  silver  deposited  as  the  chemical 
equivalent  of  copper  is  to  the  chemical  equivalent  of  silver. 
But  it  is  found  that  the  same  difficulty  exists  as  with  finding 
the  chemical  equivalents ;  namely,  that,  if  there  are  two  or 
more  different  compounds  between  the  same  elements,  then 
as  many  different  electrical  equivalents  are  obtained.  Thus 
copper  gives  31.6  and  63.2,  and  mercury  gives  99.9  and 
199.8.  So  we  have  still  the  same  question  as  to  which  num- 
ber shall  be  taken  as  the  atomic  weight. 

EXPERIMENT    14 
Electrical  Equivalents 

Have  ready  two  Daniel  cells.  In  a  beaker,  put  a  solution 
of  copper  sulfate,  and  in  another,  put  a  solution  of  nickel 


LAWS    AND    THEORIES    OF    CHEMISTRY  113 

sulfate.  Cut  out  two  copper  plates  three  centimeters  wide 
by  six  long,  making  a  hole  in  each,  and  weigh  one  care- 
fully. Cut  out  two  similar 
nickel  plates  of  the  same 
size,  and  weigh  one.  Place 
the  copper  plates  in  the 
copper  sulfate,  and  the 
nickel  plates  in  the  nickel 
sulfate.  Connect  the  cop- 
per plate  that  was  not 
weighed  with  the  weighed  nickel  plate  by  means  of  a  wire, 
and  then  put  the  two  solutions  in  circuit  with  the  two 
Daniel  cells,  so  that  the  weighed  plates  are  the  cathodes. 
Allow  the  current  to  pass  for  about  twenty  minutes  ;  then 
disconnect,  dry,  and  weigh  the  plates  again. 
Arrange  your  calculations  as  follows :  — 

Wt.  cathode  copper  plat3    : 

Wt.  after 

Difference 

Wt.  cathode  nickel  plate    = 

Wt.  after 

Difference 

Compare  the  weights  of  the  metals  deposited,  and  see 
whether  they  are  in  the  same  ratio  as  the  chemical  equiva- 
lents of  the  two  metals,  63.2  and  58.6. 

Specific  Heat.  We  have  noticed  in  our  every  day  expe- 
rience that  it  requires  a  different  amount  of  heat  to  raise  the 
temperature  of  various  substances  the  same  number  of 
degrees.  Since  heat  is  a  form  of  energy,  it  can  be  measured. 
To  measure  a  quantity  of  heat,  it  will  be  necessary  to  use 
some  definite  quantity  as  a  unit  of  measurement.  The  unit 
used  is  the  calorie.  A  calorie  is  that  quantity  of  heat  that  is 
used  up  in  raising  one  gram  of  water  one  degree  centigrade* 


114       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Therefore  to  raise  one  gram  of  water  2°,  or  two  grams   i° 
requires  two  calories.      The  amount  of  heat,  measured  in  calo- 
ries, necessary  to  raise  one  gram  of  a  substance  i°  C.,  is  called 
the  specific  heat  of  that  substance.     Below  is  given  the  specific 
heat  of  a  number  of  the  elements. 

Iron  .114  Copper  .094         Silver  .057       Gold  .032 

Nickel  .108         Zinc  .095  Tin  .055          Lead  .031 

Apparatus  used  in  Determining  Specific  Heat.  To  find  the 
specific  heat  of  elements  experimentally,  two  pieces  of 
apparatus  are  necessary  ;  first,  a  vessel,  called  a  calorimeter, 
to  hold  the  substance.  An  ordinary  nickel-plated  lemonade 
shaker  will  answer  very  well.  Secondly,  we  must  have 
some  sort  of  apparatus  in  which  the  substance  can  be  heated 
to  a  constant  temperature.  Such  an  apparatus  (see  illustra- 
tion, p.  115)  can  be  obtained  at  small  cost  from  dealers  in 
laboratory  supplies,  or  one  can  be  set  up  as  follows.  Have 
ready  a  large  beaker  of  about  1000  cc.  capacity,  one-quarter 
full  of  water.  Then  cut  a  thin  board  in  the  shape  of  a  ladle, 
large  enough  to  make  the  broad  part  cover  the  large  beaker. 
Cut  a  hole  in  the  center  of  the  broad  part  large  enough  to 
hold  another  beaker  of  about  200  cc.  capacity. 


EXPERIMENT    15 
Specific  Heat  of  Lead 

To  determine  the  specific  heat  of  lead  with  any  degree  of 
accuracy  by  means  of  the  apparatus  here  used,  we  shall 
assume  that  the  specific  heat  of  the  material  of  the  calorim- 
eter is  known.  In  the  case  of  brass  it  is  .094,  in  the  case 
of  glass  .2. 


LAWS  AND  THEORIES  OF  CHEMISTRY      I  I  $ 

Weigh  out  500  grms.  of  fine  shot,  and  place  the  shot  in 
the  small  beaker.  Set  the  small  beaker  in  the  large  beaker, 
and  heat  the  water  in  the  large  beaker  to  boiling.  Cover 
the  small  beaker  with  a  piece  of  cardboard.  Meanwhile 
weigh  the  calorimeter,  and  add  exactly  100  grms.  of  water 


cooled  about  8°  below  the  temperature  of  the  room.  Stir  the 
shot  occasionally  and  thoroughly  with  a  thermometer,  and 
note  when  the  temperature  becomes  constant.  Then  pour 
the  shot  quickly  into  the  calorimeter,  stirring  with  another 
thermometer,  and  note  the  temperature  of  the  mixture.  (The 
illustration  shows  apparatus  mentioned  on  p.  114.) 

Let  us  see  what  has  taken  place.  The  amount  of  heat 
lost  by  the  shot  (we  neglect  what  little  heat  has  been  radiated 
off)  must  be  equal  to  the  amount  of  heat  gained  by  the  water 
and  the  calorimeter. 


Il6       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Arrange  your  calculations  as  follows  :  — 

Weight  of  the  shot  =    s    = 

Weight  of  the  calorimeter  =    c    = 

Weight  of  the  water  =   w  = 

Temperature  of  the  shot  =  ts    = 

Temperature  of  the  water  =  tw  = 

Temperature  of  the  mixture  =  tm  = 
Specific  heat  of  substance  of  cal.  =    h   =  .094 

Amount  of  heat  lost  by  shot  =  amount  of  heat  gained  by 
water  -(-  amount  of  heat  gained  by  the  calorimeter. 

Let  x  =  the  specific  heat  of  lead. 

Then,  since  i  gram  of  lead  loses  x  calories  in  falling  i°, 
s  grams  of  shot  will  lose  s  (ta  —  /m)  x  calories  in  falling  4°  — 
/m°.  Since  one  gram  of  water  gains  one  calorie  in  being 
heated  one  degree,  w  grams  of  water  will  gain  w  (tm  —  tw) 
calories  in  being  raised  /w°  —  tw°.  Since  one  gram  of  the 
substance  of  the  calorimeter  gains  h  calories  in  being  raised 
one  degree,  c  grams  will  gain  eh  (tm  —  4,)  calories  in  being 
raised  tm°  —  tw°.  Therefore  we  have  the  original  statement 
expressed  algebraically. 

s  (ts  — tm)  x=w  (tm  — tw)  +  ch  (tm  — tw) 

Substituting  your  observed  values,  you  should  obtain,  by 
solving  for  x,  a  number  somewhere  near  .03 1  as  the  specific 
heat  of  lead. 

EXPERIMENT    16 
Specific  Heat  of  Iron 

In  the  same  way,  find  the  specific  heat  of  iron,  using  iron 
filings  free  from  grease  or  oil.  You  should  obtain  some- 
where near  the  number  .114. 


LAWS    AND    THEORIES    OF    CHEMISTRY 


117 


Thermal  Equivalents.  In  1819,  two  French  chemists, 
Dulong  and  Petit,  noticed  that  there  existed  a  simple  rela- 
tion between  the  chemical  equivalents  of  the  elements  and 
their  specific  heats ;  namely,  that  the  product  of  the  two 
numbers  always  approximates  6.4,  or  half  that  number.  In 
most  cases  it  is  6.4.  The  following  table  shows  this  remark- 
able relation.  The  chemical  equivalents  are  doubled  where 
it  is  necessary  to  obtain  approximately  the  product  6.4.* 


ELEMENT. 

U  H 
ll 

SPECIFIC 
HEAT. 

PRODUCT. 

ELEMENT. 

ATOMIC 
WEIGHT. 

SPECIFIC 
HEAT. 

PRODUCT. 

Iron 

55.6 

.114 

6.4 

Silver 

lOJ.II 

•057 

6.1 

Nickel 

58.24 

.108 

6-3 

Tin 

118.15 

•055 

6.5 

Copper 

63.12 

.C94 

6.0 

Gold 

195-74 

.032 

6-3 

Zinc 

64.91 

.095 

6.1 

Lead 

205.36 

.031 

6.4 

The  product  of  the  atomic  weight  and  the  specific  heat 
varies  somewhat,  of  course ;  but  it  is  only  reasonable  to  sup- 
pose that,  if  our  numbers  were  absolutely  correct,  this  product 
would  always  be  the  same.  The  quotient  of  the  number  in 
the  last  column  divided  by  the  specific  heat  is  called  the 
thermal  equivalent. 

Explanation.  Here  we  have  a  method  of  aiding  us  in 
determining  what  numbers  we  shall  take  for  the  atomic 
weights.  The  only  way  we  can  explain  the  relation  noted 
in  the  last  paragraph  is  by  saying  that  all  atoms  have  the 

*  This  product  will  probably  be  somewhat  smaller  than  6.4  as  the 
atomic  weights. are  more  and  more  accurately  determined. 


I  1 8       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

same  capacity  for  heat ;  that  is,  it  requires  the  same  amount 
of  heat  to  raise  every  atom,  no  matter  of  what  kind,  one 
degree  in  te?npe?'ature.  For  example,  the  atomic  weight  of 
lead  is  205  microcriths  (a  microcrith  being  the  weight  of  an 
atom  of  hydrogen),  and  that  of  iron  is  56  microcriths.  If 
we  assume  that  it  requires  the  same  amount  of  heat  to  raise 
an  atom  of  each  one  degree,  it  will  take  205  as  much  to 
raise  one  microcrith  of  lead  one  degree,  and  ^  as  much  to 
raise  one  microcrith  of  iron  one  degree.  But  the  atomic 
weights  must  bear  to  each  other  the  same  ratio  as  their  mass 
weights,  therefore  the  amount  of  heat  necessary  to  raise  one 
gram- of  lead  one  degree  must  be  to  the  amount  of  heat  ne- 
cessary to  raise  one  gram  of  iron  one  degree  as  ^J^  is 
to  g1^,  or  as  56  is  to  205.  The  specific  heats  of  lead  and 
iron,  .031  and  .114  respectively,  are  to  each  other  in  this 
ratio.  We  are  then  justified  in  believing  that  our  assumption 
—  that  every  atom  requires  the  same  amount  of  heat  to  raise 
it  one  degree  in  temperature  —  is  true. 

We  have  therefore  considerable  evidence  to  warrant  us  in 
believing  that 

atomic  weight  x  specific  heat  =  a  constant,  i.e.,  6.4 

From  this  we  have 

6.4 


atomic  weight  = 


specific  heat 
or 

6.4 


specific  heat  = 


atomic  weight 

From  Experiment  13,  we  find  that  32.5  grams  of  zinc  re- 
place one  gram  of  hydrogen.  Suppose  we  take  32.5  for 
the  atomic  weight,  and  divide  6.4  by  this  number.  We 
obtain  .193  for  the  specific  heat.  Now  let  us  double  the 
number  32.5,  and  divide  6.4  by  the  result.  We  then  obtain 
.096  for  the  specific  heat.  Since  the  specific  heat  of  zinc  is 


LAWS    AND    THEORIES    OF    CHEMISTRY  I  IQ 

found  by  experiment  to  be  .0955,  we  are  justified  in  believ- 
ing that  the  atomic  weight  is  65,  just  twice  its  chemical 
equivalence. 

Isomorphic  or  Crystallographic  Equivalents.  A  short  time 
after  Dulong  and  Petit  made  their  discovery,  Mitscherlich, 
a  German  chemist,  found  that  certain  elements  can  replace 
others  in  a  compound  without  changing  the  crystalline  form. 
Such  elements  are  said  to  be  isomorphous.  The  replace- 
ment always  takes  place  in  definite  quantities.  These 
quantities  are  called  the  Crystallographic  equivalents. 

The  chemical  equivalence  of  silver  is  107.  In  a  crystal, 
107  parts  by  weight  of  silver  can  be  made  indirectly  to  re- 
place 65  parts  of  zinc;  hence  we  say  that  107  parts  of  silver 
is  the  Crystallographic  equivalent  of  65  parts  of  zinc.  But 
the  chemical  equivalence  of  zinc  is  32.5.  The  only  explana- 
tion we  can  give  of  this  is  that  one  atom  of  zinc  replaces 
two  atoms  of  hydrogen;  and  hence  the  atomic  weight  of 
zinc  is  not  the  same  as  its  chemical  equivalence,  but  twice 
32.5,  or  65.  This  is  evidently  in  harmony  with  what  we 
learned  in  regard  to  the  atomic  weight  of  zinc  from  the  Law 
of  Dulong  and  Petit.  In  like  manner,  the  atomic  weights  of 
other  elements  have  been  investigated,  and  it  has  been  found 
that  this  principle  is  a  great  aid  in  determining  the  right 
number  to  use  in  many  cases.  It  has  not  however  been 
found  to  be  universally  true. 


EXPERIMENT    17 
Law  of  Definite  Proportions  by  Volume 

Refer  to  Experiment  6  <?,  Part  I.,  and  note  the  respective 
volumes  of  hydrogen  and  oxygen  that  unite  to  form  water. 
From  Experiment  24  b,  Part  I.,  note  the  number  of  volumes 


I2O       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

of  hydrogen  that  unite  with  one  volume  of  chlorin  to  form 
two  volumes  of  hydrochloric  acid  gas. 

Statement  of  the  Law.  From  such  experiments  as  the  pre- 
ceding, it  has  been  found  that,  when  gases  unite  to  form  com- 
pounds, the  relative  volumes  bear  to  each  other  a  simple  ratio. 
This  is  called  the  Law  of  Definite  Proportions  by  Volume, 
or  the  Law  of  Gay  Lussac. 

Avogadro  V  Hypothesis.  In  order  to  explain  the  above 
fact,  it  will  be  necessary  for  us  to  make  the  hypothesis  that 
equal  volumes  of  gases,  temperature  and  pressure  being  the  same, 
contain  an  equal  number  of  molecules.  This  is  called  the 
Hypothesis  of  Avogadro.  (Avogadro  was  an  Italian  physi- 
cist of  the  early  part  of  the  nineteenth  century.)  That  this 
assumption  is  true,  is  indicated  by  both  physical  and  chemi- 
cal considerations.  For  instance,  we  know  that  all  gases, 
under  changes  of  pressure,  act  in  the  same  manner ;  this  is 
also  true  for  changes  of  temperature.  Besides,  by  accepting 
the  atomic  theory  and  arguing  from  a  purely  mathematical 
standpoint,  the  truth  of  this  hypothesis  has  been  shown  to  be 
an  absolute  necessity. 

Molecular  Weights  of  Gaseous  Elements  and  Compounds. 
It  must  follow  then,  if  equal  volumes  of  gases  contain  an 
equal  number  of  molecules,  that  the  ratio  of  the  weight 
of  a  molecule  of  an  element  or  compound  to  the  weight 
of  a  molecule  of  hydrogen  is  the  same  as  the  ratio  of 
the  weights  of  equal  volumes  of  the  two  gases  ;  that  is,  the 
same  as  the  density  of  the  element  or  compound.  Thus  the 
molecular  weight  of  carbon  dioxid  compared  to  the  wreight  of 

a  molecule  of    hydrogen    must  be    — l^Z-  —  22.      It  must  be 

.09 

borne  in  mind,  however,  that  we  have  not  yet  determined 
how  many  atoms  there  are  in  the  hydrogen  molecule. 
When  we  have  done  that,  we  can  obtain  the  molecular 


LAWS    AND    THEORIES    OF    CHEMISTRY 


121 


weight  in  terms  of  the   hydrogen   atom.     This  we  proceed 
to  do. 

Number  of  Atoms  in  the  Hydrogen  Molecule.  We  shall  be 
compelled  to  believe  that  a  molecule  of  hydrogen  contains  at 
least  two  atoms.  For  we  know  by  Experiment  24  £,  Part 
I.,  that  one  volume  of  hydrogen  unites  with  one  volume 
of  chlorin  to  form  two  volumes  of  hydrochloric  acid  gas. 
Let  us  suppose  that  in  the  one  volume  of  hydrogen  there 
are  1000  molecules,  then  from  Avogadro's  Hypothesis  there 
must  be  1000  molecules  in  the  one  volume  of  chlorin,  and 
2000  molecules  in  the  two  volumes  of  hvdrochloric  acid. 


Hydrogen 


Chlorin 


Hydrochloric  Acid 


IOOO 

IOOO 

Thus : 


But  every  hydrochloric  acid  molecule  is  shown  by  analysis 
to  be  made  up  of  hydrogen  and  chlorin  ;  therefore  in  the 
2000  molecules  of  hydrochloric  acid  there  must  be  at  least 
2000  atoms  of  hydrogen.  Now  these  2000  atoms  of  hydro- 
gen were  originally  in  the  1000  molecules  of  hydrogen; 
therefore  in  these  1000  molecules  there  must  be  at  least  2000 
atoms,  and  thus  each  molecule  must  contain  at  least  two 
atoms.  In  all  the  chemical  investigations  that  have  been 
made,  nothing  has  ever  been  found  that  would  seem  to  indi- 
cate that  the  hydrogen  molecule  contains  more  than  two 
atoms ;  so  we  shall  accept  this  number.  Likewise  each 
molecule  of  chlorin  must  contain  at  least  two  atoms. 

Number  of  Atoms  in  the  Oxygen  Molecule.  It  is  proved  by 
experiment  that  two  volumes  of  hydrogen  unite  with  one  vol- 
ume of  oxygen  to  form  two  volumes  of  water  vapor.  Suppose, 
as  before,  that  one  volume  contains  1000  molecules.  We 
then  have 


122       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 
Hydrogen  Water  Vapor 


Oxygen 


Each  of  the  2000  molecules  of  water  vapor  must  contain 
at  least  one  atom  of  oxygen  ;  but  these  2000  atoms  came  from 
the  1000  molecules  of  oxygen;  hence  each  oxygen  molecule 
must  contain  at  least  two  atoms.  The  only  thing  that  makes 
us  believe  oxygen  ever  contains  more  than  two  atoms  is  the 
fact  that,  when  a  silent  electric  discharge  is  passed  through 
dry  oxygen  gas,  it  decreases  to  two-thirds  of  its  former  voi- 
ume.  This  is  easily  explained  in  the  light  of  Avogadro's 
Hypothesis.  Suppose,  as  before,  one  volume  contains  1000 
molecules. 

Then 

Oxygen  Ozone 


1000 

IOOO 

IOOO 

= 

1000 

1000 

Nothing  is  simpler  than  to  believe  that  three  oxygen  mole- 
cules, each  containing  two  atoms,  have  been  broken  up,  and 
then  reunited  to  form  two  molecules  each  containing  three 
atoms.  Hence  ozone  differs  from  oxygen  in  that  it  contains 
three  atoms  to  the  molecule  instead  of  two.  Thus  we  see 
that  we  have  a  very  rational  explanation  of  allotropic  forms. 
Atomic  Weight  of  Oxygen.  In  the  light  of  what  we  have 
seen,  it  follows  that  the  molecular  weight  of  any  gaseous 
compound  may  be  obtained  by  multiplying  its  density  by  two. 
Thus,  since  a  molecule  of  carbon  dioxid  is  22  times  as 
heavy  as  a  molecule  of  hydrogen,  and  since  a  molecule  of 


LAWS    AND    THEORIES    OF    CHEMISTRY  123 

hydrogen  contains  two  atoms,  therefore  a  molecule  of  carbon 
dioxid  will  be  44  (i.e.,  2x22)  times  as  heavy  as  an  atom  of 
hydrogen.  In  like  manner,  the  molecular  weight  of  oxygen 
is  found  to  be  32  ;  and,  since  there  are  two  atoms  in  the 
molecule,  its  atomic  weight  is  16. 

Molecular  Weights  of  Compounds  not  Capable  of  Being 
Volatilized,  To  solid  compounds  not  capable  of  being  vol- 
atilized, we  are  unable  to  apply  the  ordinary  method  of  find- 
ing the  molecular  weight.  We  are  obliged  to  attack  the 
question  in  another  way.  This  is  best  illustrated  by  the 
following  experiment. 

EXPERIMENT    18 

Molecular  Weight  of  Potassium  Chlorate 

Dry  in  an  oven  at  about  150°  a  quantity  of  powdered  c.p. 
potassium  chlorate.  Weigh  a  porcelain  crucible  and  its 
cover  on  the  delicate  balance.  Then  weigh  out  in  the  cru- 
cible about  2  grms.  of  the  dry  chlorate.  Heat  it  on  a  triangle, 
gently  at  first,  holding  the  cover  over  the  crucible  by  means 
of  a  pair  of  long  forceps  in  order  to  catch  any  of  the  melted 
substance  that  may  spatter.  Be  careful  not  to  heat  too 
strongly  at  first,  since  the  chlorate  may  bubble  over  the 
sides  of  the  crucible  and  spoil  the  experiment.  When  all 
the  oxygen  is  driven  off,  cool  the  crucible  in  a  desiccator  and 
weigh  it  again.  Continue  to  heat  and  weigh  it,  until  you 
obtain  a  constant  weight.  From  the  numbers  obtained,  it  is 
possible  to  get  the  molecular  weight. 

We  know  that  potassium  chlorate  is  composed  of  potas- 
sium, chlorin,  and  oxygen.  If  we  knew  how  many  atoms  of 
oxygen  there  are  in  the  potassium  chlorate  molecule,  we 
could  make  our  calculation  easily.  Now  chemists  have  de- 


124       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

termined  (it  is  a  problem  beyond  elementary  students)  that  a 
molecule  of  potassium  chlorate  must  contain  either  three  or 
a  multiple  of  three  atoms  of  oxygen.  Three  has  been  chosen 
as  the  more  probable  number.  If  then  there  are  three  atoms 
of  oxygen  in  the  molecule,  the  total  weight  of  the  oxygen 
atoms  will  be  48  (i.e.,  3  X  16)  microcriths.  Let  x  be  the 
molecular  weight  of  potassium  chlorate,  and  make  the  pro- 
portion x  :  48  ::  wt.  of  potassium  chlorate  taken  :  wt.  of  oxygen 
lost.  Solving  the  proportion  for  x,  you  should  obtain  some- 
where near  the  number  122. 

EXPERIMENT    19 
Molecular  Weight  of  Potassium  Chlorid 

Knowing  the  molecular  weight  of  potassium  chlorate,  it  is 
easy  for  us  to  find  the  molecular  weight  of  potassium  chlorid. 
From  Experiment  6  and  25  c,  Part  I.,  we  know  that  the  resi- 
due left  in  the  crucible  after  heating  the  chlorate  is  potassium 
chlorid ;  hence  the  molecular  weight  of  potassium  chlorate 
minus  48  is  the  molecular  weight  of  potassium  chlorid. 

By  similar  methods  the  molecular  weights  of  a  large  num- 
ber of  compounds  that  cannot  be  volatilized  have  been  deter- 
mined. But  it  must  be  remembered  that  we  cannot  be  sure 
that  the  numbers  we  obtain  by  the  chemical  method  are  the 
correct  ones,  until  we  have  studied  a  great  number  of  chemi- 
cal changes  into  which  the  elements  that  constitute  the  com- 
pounds enter.  Besides  this,  we  must  verify  these  numbers 
by  means  of  other  facts. 

Atomic  Weights  in  General.  By  comparing  all  the  various 
numbers  obtained  from  experiments  for  determining  chemi- 
cal, electrical,  isomorphic,  and  thermal  equivalents,  and 
molecular  weights  obtained  by  the  physical  and  chemical 


LAWS    AND    THEORIES    OF    CHEMISTRY  12$ 

methods,  chemists  have  obtained  a  series  of  numbers  which 
represent  so  accurately  the  atomic  weights  of  the  elements 
that  there  is  practically  no  doubt  of  their  truth  within  very 
small  prescribed  limits.  These  numbers  are  used  without 
showing  appreciable  error  in  numberless  chemical  investiga- 
tions, and  they  are  to-day  accepted  to  be  as  near  the  actual 
numbers  as  our  present  methods  of  analysis  admit. 

Periodicity  of  the  Elements.  If  we  write  horizontally  in 
order,  in  seven .  columns,  the  numbers  representing  the 
atomic  weights  of  the  elements  (omitting  hydrogen),  we  shall 
see  that  those  in  the  same  vertical  column  resemble  each 
other  to  a  remarkable  degree  both  in  their  chemical  proper- 
ties and  in  the  similarity  of  their  compounds.  A  short  study 
of  this  table  will  be  profitable. 

In  the  first  column,  we  find  first  the  strong  alkalis ;  in  the 
second,  the  closely  allied  elements  calcium,  barium,  and 
strontium,  and  also  the  natural  group  magnesium,  zinc,  and 
cadmium.  In  the  third  are  found  most  of  the  rarer  ele- 
ments ;  in  the  fifth,  we  have  the  similar  elements  oxygen, 
sulfur,  selenium,  and  tellurium;  and  in  the  seventh,  we  have 
the  halogens.  Thus  it  seems  that  the  properties  of  the  ele- 
ments depend  upon  their  atomic  weights,  with  similar  prop- 
erties recurring  at  regular  intervals.  Besides  the  regular 
column,  an  extra  column  is  added  for  those  elements  which 
for  some  unknown  reason  do  not  seem  to  fit  in  with  the  regu- 
lar groups.  It  will  be  noticed  that,  in  the  irregular  groups, 
the  elements  have  almost  the  same  atomic  weights  and  simi- 
lar properties.  Vacant  spaces  have  been  left,  as  there  are 
at  present  no  elements  known  whose  atomic  weights  fit  in 
those  spaces.  When  the  table  was  first  made  out,  Mendelejeff, 
to  whom  we  are  indebted  for  this  remarkable  discovery,  left 
two  vacant  spaces  between  the  elements  zinc  and  arsenic. 
Mendelejeff  predicted  the  properties  of  both  these  elements, 


126       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


to ~ 
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LAWS    AND    THEORIES    OF    CHEMISTRY 

which  have  since  been  discovered  and  found  to  agree  in  a 
wonderful  degree  with  what  had  been  expected. 

This  table  has  been  of  service  in  determining  what  numbers 
should  be  taken  as  the  atomic  weights  of  certain  elements. 
For  instance,  the  atomic  weight  of  uranium  was  supposed  to 
be  120,  but  that  number  threw  it  out  of  place  in  the  table. 
The  number  was  then  doubled,  which  placed  it  in  the  posi- 
tion in  the  table  where  from  its  properties  it  evidently  be- 
longed. This  weight  has  also  been  verified  by  its  specific 
heat  determination. 

Proufs  Hypothesis.  In  1815,  an  English  chemist,  Prout, 
advanced  the  theory  that,  inasmuch  as  the  atomic  weights  of 
the  elements  were  simple  multiples  of  the  atomic  weight  of 
hydrogen,  therefore  they  were  all  composed  of  definite  por- 
tions of  one  kind  of  matter,  and  that  the  quantity  of  this 
kind  of  matter  in  an  atom  determined  its  properties.  How- 
ever attractive  this  theory  may  be,  many  of  the  greatest  in- 
vestigators believe  it  to  be  untenable  from  the  fact  that,  as 
accuracy  in  the  determination  of  atomic  weights  has  in- 
creased, it  has  been  found  that  comparatively  few  of  them 
can  be  expressed  by  whole  numbers. 

Symbols.  We  are  now  ready  for  some  simple  method  of 
representing  chemical  changes,  not  only  qualitatively  as  in 
Part  I.,  but  quantitatively  also.  Chemists  have  agreed  to 
represent  the  atom  of  each  element  by  a  letter,  which  shall 
stand  both  for  the  atom  itself  and  also  for  its  atomic  weight. 
Thus  the  letter  H  stands  for  an  atom  of  hydrogen,  and  also 
for  its  atomic  weight  i  ;  the  letter  O  stands  for  an  atom  of 
oxygen,  and  also  for  its  atomic  weight  16.  By  combining 
these  letters,  we  may  represent  molecules.  Thus  H2SO4 
signifies  a  molecule  of  sulfuric  acid,  molecular  weight  98  ; 
the  small  numerals  mean  that  there  are  two  atoms  of  hydrogen 
and  four  atoms  of  oxygen,  besides  the  single  atom  of  sulfur. 


128       AN     ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


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LAWS    AND    THEORIES    OF    CHEMISTRY 

The  table  on  page  128  gives  a  list  of  the  elements  with 
their  symbols  and  atomic  weights.* 

Determination  of  Molecular  Formulae.  How  do  we  know 
that  sulfuric  acid  is  H2SO4,  that  copper  sulfate  is  CuSO4, 
that  alcohol  is  C2H6O,  etc.  ? 

First  we  must  determine  by  analysis  the  percentage  com- 
position of  the  compound.  Let  us  take  as  an  example  the 
last  one  above  mentioned,  alcohol.  When  analyzed,  alcohol 

is  found  to  contain 

Carbon  52.17% 

Hydrogen  13.05  % 

Oxygen  34 .78  % 

IOO.OO  % 

We  find  that  the  ratios  52.17:13.05:34.78  can  be  ex- 
pressed in  any  of  the  following  ways  :  — 


4 

I 

2.66 

8 

2  : 

5-32 

12 

3 

7.98 

16 

4 

10.64 

20 

5 

13-30 

24 

6 

15.96 

and 

the  last  of  which  is  practically  24:6:16.  Therefore  24 
grams  of  carbon,  6  grams  of  hydrogen,  and  16  grams  of 
oxygen  unite,  forming  46  grams  of  alcohol.  Or,  expressing 
the  quantity  of  each  element  in  microcriths,  24  microcriths 
of  carbon,  6  microcriths  of  hydrogen,  and  16  microcriths  of 
oxygen  unite,  forming  46  microcriths  of  alcohol.  But,  since 
an  atom  of  carbon  weighs  12  microcriths,  an  atom  of  hydro- 
gen i  microcrith,  and  an  atom  of  oxygen  16  microcriths,  we 
see  at  once  that  the  simplest  composition  of  a  molecule  of 
alcohol  would  be  two  atoms  of  carbon,  six  atoms  of  hydrogen, 
and  one  atom  of  oxygen,  i.e.,  C2H6O. 

*  From  report  of  American  Chemical  Society,  1899. 


I3O       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

We  might  just  as  well  use  the  ratios 

48  :  12  :  32 
Or         64  :  18  :  48 

We  should  then  have  obtained  either  the  formula  C4H12O2  or 
C6H18O3.  The  only  way  we  have  of  deciding  which  is  cor- 
rect is  by  comparing  the  molecular  weights  which  these  for- 
mulae give,  with  the  density  of  alcohol  in  the  state  of  vapor. 

C2  H6  0  gives  a  molecular  weight  of  24  +  6  +  16  =  46 
C4  Hi2  02  gives  a  molecular  weight  of  48  +  12  +  32  =  92 
Ce  HIS  Oa  gives  a  molecular  weight  of  72  +  18  +  48  =  138 

By  weighing  alcohol  in  the  state  of  vapor  and  comparing  it 
with  hydrogen,  we  found  in  Experiment  12,  Part  II.  that  its 
destiny  is  23.  Therefore  its  molecular  weight  must  be 
2  x  23,  or  46.  Hence  we  take  the  formula  C2H6O  as  the 
correct  one.  In  the  case  of  compounds  whose  density  can- 
not be  found,  we  take  the  simplest  formula. 

The  simplest  way  numerically  to  find  the  formula  when 
the  percentage  composition  is  given  is  as  follows : 

Carbon  52.17 
Hydrogen  13.05 
Oxygen  34.78 

Rule.  Divide  the  percentage  of  each  element  by  its  atomic 
weight.  Find  what  the  resulting  ratios  become,  if  the  smallest 
number  be  taken  as  unity.  If  there  is  only  one  atom  of  that 
element,  the  ratios  will  then  be  expressed  by  whole  numbers.  If 
the  resulting  ratios  are  not  whole  numbers,  try  successively  what 
the  ratios  would  become,  if  the  smallest  number  were  taken  as 
2,  j,  etc. 

Examples.  Deduce  the  formulae  for  the  following  sub- 
stances :  — 


LAWS    AND    THEORIES    OF    CHEMISTRY 


I. 

2_ 

S- 
6. 

f 

r 

I 
r 

I 

r 

1 

r 
j 
^ 

L 
1 

Hydrogen 
Oxygen 

Density,  9 

Hydrogen 
Oxygen 

II.  12 

88.88 

IOO.OO 

5-88 
94.12 

IOO.OO 

Manganese  72.05 
Oxygen        27.95 

Silicon 
Oxygen 

Iron 
Oxygen 

Carbon 
Hydrogen 

IOO.OO 

46.67 

53-33 

IOO.OO 

70.01 

29-99 

IOO.OO 

92.3 

7-7 

100.00 

Density,  13. 

f  Copper 
1  Sulfur 
7*1  Oxygen 

f  Hydrogen 
~  J  Phosphorus 
j  Oxygen 

f  Potassium 
J  Nitrogen 
j  Oxygen 

f  Nitrogen 
|  Hydrogen 
10.  -|  Carbon 
Oxygen 

39-62 
20.13 
40.25 

IOO.OO 

3.06 

31.64 
65-30 

IOO.OO 

45-95 
16.45 

37-6o 

IOO.OO 

29.17 

8.33 
12.50 
50.00 

100.00 

Valence.     Let  us  suppose  that  we  know  the  molecular  for- 
mulae for  all  the  known  compounds. 

Then,  if  we  arrange  a  few  of  the  simplest  thus :  — 

I.  II.  III.  IV. 

HF  H20  H3N  H4C 

H  Cl  Ho  S  H3  P  Hi  Si 

H  Br  H2  Se  H3  As 

HI  H2Te  H3Sb  

we  shall   see  that  an   atom  of  each  element  possesses  the 
power  of  uniting  with  a  certain  number  of  hydrogen  atoms. 
Again,  let  us  make  a  table  thus :  — 
Acids  Salts 


HC1 

AgCl 

CuCl2 

A1C18 

HN03 

AgN03 

Cu  (N  03)2 

Al  (N  Os)3 

H2S  04 

Ag2  S  04 

CuS04 

A12  (S  04)3 

H3P04 

Ag3  P  04 

Cu3  (P  04)2 

A1P04 

132         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

We  here  see  that  an  atom  of  each  element  possesses  the 
power  of  replacing  a  certain  number  of  hydrogen  atoms  in  a 
compound.  This  property  of  an  atom  is  called  its  valence. 
We  might  define  valence  as  the  quantity  of  combining  power  an 
atom  has  compared  with  that  of  an  atom  of  hydrogen.  Atoms 
that  combine  with  or  replace  an  atom  of  hydrogen  are  called 
monovalent,  or  monads.  Those  which  combine  with  or 
replace  two  atoms  of  hydrogen  are  called  divalent,  or  dyads. 
There  are  also  trivalent,  tetravalent,  etc.,  atoms.  For  some 
of  the  elements,  the  number  representing  the  valence  varies. 
For  instance,  nitrogen  has  a  variable  valence,  being  some- 
times triad  and  sometimes  pentad. 

The  following  table  shows  the  valence  of  the  more  common 
elements. 

Na  i 
Ni  2 

0       2 

P     3,5 
Pt    4 
Pb  2 

S     2,4,6 
Si   4 
Sn  2,4 
Sr  2 
Sb  3,5 
Zn  2 

Graphic  Formulae.  Although  we  do  not  know  the  nature 
of  chemical  affinity,  still,  as  we  have  seen,  we  can  express  it 
quantitatively.  We  can  represent  monad,  dyad,  triad,  etc., 
valence  by  lines ;  a  monad  element  being  written  H  -  ,  a 

I 
dyad,  -  O  -  ,  a  triad,  B  - ,  and  so  on. 

Let  us  start  with  the  compound  H2O,  and  write  it  H  -  O-  H, 


Al   3 

Cr  3 

As  3,  5 

CU  2 

Agi 

F    i 

B    3 

Fe  2,4 

Ba2 

H    i 

Bi  3,  5 

Hg2 

Br  i  (3,  5,  7) 

I      i(5) 

C    4(2) 

K    i 

Ca  2 

Li    i 

Cl   i  (3,  5,  7) 

Mg2 

Cd2 

Mn2 

Co  2 

N    3,5 

LAWS    AND    THEORIES    OF    CHEMISTRY  133 

TT 

or  _          O,  understanding  by  this  that  one  atom  of  oxygen 
rl   ^ 

with  a  valence  of  two  is  united  with  two  atoms  of  hydrogen, 
each  having  a  valence  of  one.  Suppose  now  that  we  replace 
the  monad  hydrogen  atoms  by  two  monad  sodium  atoms. 
We  shall  then  have  for  sodium  oxid  the  formula 

Na  — 0  — Na  or  JJ*  >  0 

In  like  manner  for  potassium  oxid,  we  should  have 
K  —  O  —  K.  If  we  replace  the  two  monad  hydrogen  atoms 
by  a  single  dyad  atom,  such  as  calcium,  we  should  have 
Ca  =  O  for  calcium  oxid.  Using  the  triad  aluminum  atom, 
we  should  have  to  replace  the  six  monad  hydrogen  atoms 
of  three  molecules  of  water  by  two  aluminum  atoms. 

H-  o 

H-  ..^0 


H-" 

In  like  manner  other  graphic  formulae  of  other  oxids  may 
be  written. 

The  corresponding  hydroxids,  of  course,  will  be 

n      „         /O— H 
Na  — 0  — H,  K  — 0  — H,  Ca<5~JJ,  Al  —  0  —  H 

the  metals  replacing  only  one  hydrogen  atom  from  each  water 
molecule. 

In    the    case   of    an    acid,  we  write    H  —  Cl    for    hydro- 
chloric, H  —  Br  for  hydrobromic,  etc.     We  then  have : 


SODIUM  SALTS 

CALCIUM  SALTS 

ALUMINUM  SALTS 

/ 

Cl 

Na  —  Cl 

Ca<Cl 

Al  — 

Cl 

Na  —  Br,  etc. 

Ca<Br,etc. 

/ 

Al  — 

Br 
Br 

^ 

Br,  etc. 

134      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

We  have  seen  that  all  acids  contain  hydrogen,  but  not 
necessarily  oxygen.  The  greater  number  of  acids,  however, 
contain  oxygen,  and  it  has  been  shown  by  a  number  of  in- 
vestigators that  at  least  part  of  the  oxygen  in  these  acid 
molecules  is  closely  associated  with  the  replaceable  hydro- 
gen. Thus  in  nitric,  HNO3,  we  have  the  graphic  formula 

H  —  O  —  N  J  ~.     This  group,  —  O  —  H,  which  appears  also 

in  the  bases  Na  —  O  —  H,  K  -  O  —  H,  etc.,  is  called  the 
hydroxyl  group,  and  is  evidently  a  monad.  We  know  that 
sodium  nitrate  differs  from  nitric  acid,  in  that  it  contains  an 
atom  of  sodium  in  place  of  an  atom  of  hydrogen. 

We  have,  therefore,  the  following  formulae : 

SODIUM  NITRATE  CALCIUM  NITRATE  ALUMINUM  NITRATE 

™^0 


0 


,0  — N 


Na  — 0  — N^X      Ca  X         Al  — 0  — N 


0 


In  the  same  way,  it  is  believed,  sulfuric  acid,  H2SO4,  con- 
tains two  hydroxyl  groups,  thus  :  — 

H—  0  0 


From  this  we  may  obtain  the  salts 

Na  —  0  0 

^S  ^      ,  sodium  sulfate 
Na  —  0^     ^  0 


0  0 

^  S  ^     ,  calcium  sulfate 
^     ^0 


LAWS    AND    THEORIES    OF    CHEMISTRY  135 

/o^c^o 

Al—  0^b^0 

~  ^;  S  ^  f)  >  aluminum  sulfate 

Al—  0\Q^0 
^0/S=-0 

Similarly  phosphoric  acid,  H8PO4,  contains  three  hydroxyl 
groups,  thus  :  — 

H  —  0-, 

H  —  0—  P  —  0 

H  —  0^ 
from  which  we  have  the  salts, 

Na  —  0^ 

Na  —  0—  P  —  0,  sodium  phosphate 

Na  —  O-' 


Ca  <     ^  »   calcium  phosphate 

0—  P—  0 


Al  —  0  — P  —  0,  aluminum  phosphate 

Acids  that  contain  one  hydroxyl  group  are  called  mono- 
basic, those  containing  two  such  groups  are  called  dibasic, 
those  containing  three  are  called  tribasic,  and  so  on.  When 
all  the  hydrogen  of  the  hydroxyl  groups  is  replaced  by  a 
metal,  the  resulting  salt  is  said  to  be  normal.  Evidently  it 
is  possible  that  only  the  hydrogen  of  one  hydroxyl  group  in 
a  dibasic  or  higher  acid  may  be  replaced,  thus  giving  rise  to 
acid  salts.  Hence  two  different  salts  containing  the  same 
metal  may  be  obtained  from  a  dibasic  acid,  three  from  a  tri- 
basic acid,  and  so  on.  We  may  illustrate  this  as  follows : — 

Q    f  Na  -0         ^0 

S  ^     ,  acid  sodium  sulfate 
H     —  O^    ^  0 

%  u  ^ 

Na  -  0         ^  0 

"^  S  ^     ,  normal  sodium  sulfate 


136      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


Na  —  0  ^ 

H    —  0  —  P  —  0,  dihydrogen  sodium  phosphate 
H     -0  ^ 


Na  — 0  ^ 

Na  —  0  —  P  —  0,  hydrogen  disodium  phosphate 
H    —0  ^ 

Na  — 0  ^ 

Na  —  0  —  P  —  0,  normal  sodium  phosphate 

Na  —  0  ^ 


At  present  little  is  known  concerning  the  molecular  struc- 
ture of  basic  salts.  Some  think  that  to  form  these  salts, 
additional  molecules  of  the  base  replace  part  of  the  oxygen 
of  the  acid  ;  while  others  think  that  they  are  combinations 
of  the  normal  salt  molecules  with  extra  base  molecules. 

^H 

The  molecular  formula  for  ammonia  is  NH3,  or  N  —  H 

^H 

We  have  formed  ammonium  chlorid,  NH4C1,  synthetically  by 
the  union  of  ammonia  and  hydrochloric  acid,  giving  of  course 

^H 
Cl-N   < 


and  showing  the  pentad  valence   of  nitrogen.     The   group 
NH4  is  evidently  a  monad  group,  and  as  such  appears  in 
H  ^ 

TJ 

„    ^  N  —  0  —  H,  ammonium  hydroxid 
11  ^ 

H  ^ 
H 


_  0  —    Nj    ,  ammonium  nitrate 


Q,  ammonium  sulfate 


LAWS    AND    THEORIES    OF    CHEMISTRY  137 

This  group  is  called  the  ammonium  group,  and  acts  very 
much  like  a  metal. 

Thus  the  student  will  see  the  possibility  of  representing 
through  the  eye  relations  that  would  otherwise  be  incom- 
pletely realized.  It  must  not,  however,  be  thought  that 
these  structural  formulae  are  intended  to  represent  the 
actual  position  of  the  atoms  of  a  molecule  with  respect  to 
each  other.  The  intention  is  merely  to  emphasize  certain 
relations  that  experiment  and  reason  have  shown  must  exist 
between  the  atoms  in  the  molecule. 

Examples.  Write  the  graphic  formulae  for  silver  nitrate, 
aluminum  hydroxid,  calcium  carbonate,  and  ammonium  phos- 
phate. 

Positive  and  Negative  Elements.  Chemical  compounds  can 
be  decomposed  by  electricity.  In  all  such  decompositions, 
part  of  the  atoms  appear  at  the  negative  pole,  and  part  at 
the  positive  pole.  Those  which  appear  at  the  negative  pole 
are  called  electro-positive,  or  simply  positive,  elements ;  and 
those  which  appear  at  the  positive  pole  are  called  electro- 
negative, or  simply  negative.  One  element,  however,  need 
not  always  appear  at  the  same  pole.  When  liberated  from 
its  union  with  one  element,  it  may  appear  at  the  negative 
pole ;  while,  when  liberated  from  another,  it  may  appear  at 
the  positive  pole.  The  non-metals  are  negative  with  re- 
spect to  the  metals,  but  it  is  readily  seen  that,  compared 
with  one  another,  they  may  be  one  or  the  other  according 
to  what  elements  they  are.  The  strength  of  chemical  com- 
bination depends  upon  this,  the  rule  being  that  the  more 
electrically  remote  the  elements  are,  the  stronger  their 
union. 

Naming  of  Compounds.  Binary  compounds  are  those 
composed  of  two  elements.  All  others  are  called  ternary. 
The  names  of  binary  compounds,  the  names  of  ternary  com- 


138       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

pounds  that  do  not  contain  oxygen,  and  the  names  of  com- 
pounds of  positive  elements  with  hydrogen  and  oxygen,  end  in 
id.  Examples  respectively,  NaCl,  sodium  chlorid ;  KCN, 
potassium  cyanid ;  Zn(OH)2,  zinc  hydroxid.  In  all  cases 
the  positive  element  is  named  first ;  then  the  negative  ele- 
ment or  radical,  with  the  suffix.  When  more  than  one 
compound  is  formed  between  the  same  elements,  the  name 
of  the  positive  element  ends  in  ous  or  if,  according  as  there 
is  less  or  more  of  the  negative  element  or  radical.  Examples 
Hgl,  mercur0/w  iodid;  HgI2,  mercur/r  iodid. 

An  acid  takes  its  name  from  the  characteristic  element  in 
it ;  and  when  there  is  more  than  one  acid  from  the  same 
element,  the  name  ends  in  ous  or  ic  according  as  there  is  less 
or  more  oxygen.  Examples,  H2SO3,  sulfuiw/j  acid ;  H2SO4, 
sulfur/V  acid.  If  there  are  more  than  two,  the  prefixes  hypo, 
meta,  per,  and  pyro,  are  used  also.  Examples  H2SO2,  hypo- 
sulfurous  acid ;  HC1O4,  perchloric  acid.  The  names  of 
acids  that  contain  no  oxygen  have  the  prefix  hydro. 
Example,  HC1,  hydrochloric  acid. 

Salts  take  their  names  from  the  name  of  the  positive  ele- 
ment and  that  of  the  acid.  The  positive  element  is  named 
first,  and  then  the  acid  ending  in  ite  or  ate.  The  names  of 
salts  from  "ous"  acids  end  in  ite\  those  from  "if"  acids  end 
in  ate.  Examples,  Na2SO3  sodium  sulf//!?/  N^SO^  sodium 
sultafe.  When  there  is  more  than  one  salt  from  the  same 
element  and  acid,  the  name  of  the  positive  element  ends  in 
ous  or  if,  according  as  there  is  less  or  more  of  the  acid  radi- 
cal. Examples,  HgNO3  mercur0«j  nitrate  ;  Hg(NO3)2,  mer- 
curif  nitrate. 

Writing  of  Reactions.  In  order  to  be  able  to  write  reac- 
tions by  means  of  chemical  symbols,  it  will  be  necessary 
for  the  student  to  learn  the  valence  of  each  of  the  ele- 
ments. 


LAWS    AND    THEORIES    OF    CHEMISTRY  139 

Suppose  we  wish  to  express  qualitatively  the  fact  that 
the  action  of  silver  nitrate  upon  sodium  chlorid  gives  silver 
chlorid  and  sodium  nitrate,  we  may  do  so  with  a  formula 
like  that  used  in  Part  I. 

Silver  Sodium 

Nitrogen  +         ™"    =  Nitrogen    + 


Oxygen  Oxygen  Chlorin 

In  order  to  express  it  quantitatively,  we  must  make  use  of 
the  molecular  formulae  of  the  molecules  that  act  upon  each 
other,  keeping  in  mind  the  valence  of  the  atoms  of  which 
these  molecules  are  composed.  Thus, 

Ag  N  03  +  Na  Cl  =  Na  N  03  +  Ag  Cl 

Sometimes  it  will  be  necessary  to  use  more  than  one  mole- 
cule of  each  substance,  as  in  the  case  of  the  action  of  steam 
upon  red-hot  iron. 

3  Fe  +  4  H2  0  =  Fe3  04  +  4  H2 

These  formulae  not  only  show  what  action  has  taken 
place  between  the  molecules,  but  also  how  much  matter  is 
involved.  Thus  they  indicate  the  fact  that  there  is  just  as 
much  matter  after  the  action  as  there  was  before.  For  in- 
stance, the  formula, 

Ag  N  03  +  Na  Cl  =  Na  N  Cs  +  Ag  Cl 

stands  for  the  following  sentence.  One  molecule  of  silver 
nitrate,  molecular  wt.  168.68,  acts  with  one  molecule  of 
sodium  chlorid,  molecular  wt.  58.06,  producing'one  molecule 
of  sodium  nitrate,  molecular  wt.  84.45,  and  one  molecule  of 
silver  chlorid,  molecular  wt.  142.29. 

The  student  will  be  aided  considerably  in  writing  reac- 
tions by  the  following  illustration.  Suppose  we  wish  to 
write  the  reaction  for  calcium  nitrate  and  sodium  phos- 
phate. Write  the  formulae  for  the  two  compounds  thus  :  — 

^N  Oz          Na  ^ 
Ca  +    Na  —  P  0 

Na  ^ 


I4O       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


Knowing  that  calcium  is  dyad  in  valence,  and  that  phos- 
phoric acid  is  a  tri-basic  acid,  we  see  that  we  must  take 
enough  molecules  of  calcium  nitrate  to  make  the  total  va- 
lence of  the  calcium  atoms  equal  to  6,  and  enough  sodium 
phosphate  molecules  to  make  the  total  valence  of  the 
sodium  atoms  equal  to  6,  i.e.,  in  both  cases  the  least 
common  multiple  of  2  and  3,  the  valence  of  calcium  and  the 
basicity  of  phosphoric  acid  respectively.  We  then  have 

Na 


Ca 


Ca 


Ca 


N03 
N03 
N03 
N03 
N03 
N03 


NaP04 

Na 

Na 

NaP04 

Na 


The  three  Ca  atoms  will  now  replace  the  six  sodium 
atoms,  unite  with  the  two  PO4  groups,  and  give  one  mole- 
cule of  calcium  phosphate,  indicated  thus :  — • 


N  03 

Na 

Ca 

N  03 

Na 

P04 

Na 

N  O3 

Ca 

N  03 

Na 

N08 

Na 

P04 

Ca 

N  03 

Na 

LAWS    AND    THEORIES    OF    CHEMISTRY 


The  remaining  Na's    and  NO3's    will   unite,  forming  six 
molecules  of  sodium  nitrate,  thus  :  — 


We  then  have  the  complete  equation 

3  Ca  (N  03)2  +  2  Na3  P  04  =  Ca3  (P  04)2  +  6  Na  N  Oa 


EXPERIMENT    20 

Write  the  reactions  for  all  the  experiments  in  Part  I.,  using 
chemical  symbols. 

Stoichiometry.  Since  chemical  action  takes  place  between 
molecules  of  elements  or  compounds,  in  order  to  find  the 
mass  weights  of  elements  or  compounds  formed,  knowing 
the  mass  weights  of  elements  or  compounds  used,  we  have 
only  to  make  use  of  the  following :  —  First  write  the  reaction 
representing  the  chemical  change.  Then  make  the  proportion, 
—  molecular  weight  of  the  given  substance  is  to  the  molecular 
weight  of  the  required  substance  as  the  mass  weight  of  the  given 
substance  is  to  the  mass  weight  of  the  required  substance. 


142        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Let  us  take  an  example. 

How  many  grams  of  hydrogen  will  be  evolved,  when  10 
grams  of  zinc  are  dissolved  in  sulfuric  acid  ? 

H2  S  04  +  Zn  =  Zn  S  04  +  H2 

98      +  65  =      161      +  2 

In  this  case  we  have 

65  :  2  :  :  10  :  X.      .'.  X  =  .30+  grms. 

Suppose  we  wish  to  know  the  number  of  grams  of  zinc 
sulfate  formed.  We  then  have 

65  :  161  : :  10  :  X.      .'.  X  =  24.7+  grms. 

Suppose  we  wish  to  know  the  number  of  grams  of  sulfuric 
acid  required  to  dissolve  10  grams  of  zinc.  We  then  have 

6s:g8::io:x,     .'.  x  =  15.0+  grms. 

In  case  we  wish  to  know  what  the  volume  of  the  hydrogen 
evolved  would  be  at  o°  and  760  mm.,  all  we  have  to  do  is 
to  divide  the  weight  of  the  hydrogen  by  .0896,  the  weight  of 
a  liter  of  hydrogen  at  o°  and  760  mm.  (See  Exp.  10, 
Part  II.) 

If  we  wish  to  know  what  this  volume  would  be  at  any 
required  temperature  and  pressure,  we  can  easily  find  it  by 
using  the  formula 

VP 


273+ 


(See  Exp.  6  and  7,  Part  II.) 


Examples,  i.  How  many  grams  of  potassium  chlorate  must 
be  used  to  obtain  100  grams  of  oxygen  ?  Ans.  2  64+ grms. 

2.  In  order  to  fill  a  balloon,  150  kilograms  of  hydrogen 
are  necessary.  How  much  zinc  and  sulfuric  acid  will  be 

required  to  produce  the  gas  ? 

Zinc,  4871;  kilos. 
Ans.  „.,..,  ,  ., 

Sulfuric  acid,  7350  kilos. 


LAWS  AND  THEORIES  OF  CHEMISTRY  143 

3.  What  would  the  volume  of  the  gas  be  at  o°  and  760  mm. 
pressure  ?    What  would  it  be  on  a  day  when  the  temperature 
was  20°  and  the  pressure  755  mm.  ? 

1666.6+  liters. 
Ans. 

1800.5+  liters. 

4.  How  many  grams  of  silver  nitrate  would  be  required 
to  make  20  grams  of  silver  chlorid  ?  Ans.  23.7  +  grms. 

5.  How  many  grams  of  iron  sulfid  must  be  used  to  pro- 
duce, at  o°  and  760  mm.,  10  liters  of  hydrogen  sulfid,  i  liter 
of  hydrogen  sulfid  weighing  1.52  grams?    Ans.  39.3+  grms. 

6.  Calculate  the  amount  of  manganese  dioxid  that  must 
be  used  to  produce,  at  o°  and  760  mm.,  10  liters  of  chlorin 
from  hydrochloric  acid,  i  liter  of  chlorin  weighing  3.17  grms. 

Ans.  38.9  grms. 

7.  What  weight  of  copper  would  be  used  in  making  20 
grams  of  copper  nitrate  by  dissolving  the  copper  in  nitric 
acid  ?  Ans.  6.9  grms. 

The  reactions  also  indicate  the  relative  volumes  involved 
in  the  case  of  gaseous  factors  and  products.  By  the  equation 

H2  +  C12  =  2  H  Cl 

we  indicate  that  i  volume  of  hydrogen  combines  with  i  vol- 
ume of  chlorin,  forming  2  volumes  of  hydrochloric  acid.  In 
the  same  way 

aH2  +  02  =  2  H2  0 

indicates  that  2  volumes  of  hydrogen  unite  with  i  volume  of 
oxygen,  forming  2  volumes  of  water  vapor. 

Examples,  i.  18  cc.  of  hydrogen  are  mixed  with  10  cc. 
of  chlorin  and  exploded.  What  gases  are  formed,  and  what  are 
their  volumes  ? 

2.  If  2  volumes  of  nitric  oxid  and  5  volumes  of  hydrogen 
are  united,  what  volume  of  ammonia  is  produced  ? 

3.  To  a  certain  volume  of  hydrogen  sulfid  gas,  was  added 


144     AN  ELEMENTARY   EXPERIMENTAL   CHEMISTRY 

125  cc.  of  chlorin  gas,  which  was  entirely  consumed.  What 
volume  of  hydrochloric  acid  gas  was  evolved  ? 

4.  If  150  liters  of  marsh  gas  be  exploded  with  300  liters 
of  oxygen,  what  volume  of  carbon  dioxid  would  result  ? 

Calculation  of  Percentage  Composition,  having  given  the 
Molecular  Formula.  It  is  often  required  that  the  student 
should  be  able  to  calculate  the  percentage  composition  of  a 
compound  when  the  molecular  formula  is  known.  Suppose 
we  have  given  the  formula  for  alcohol,  which  is  C2H6O. 

Carbon  ax  12  =  24 

Hydrogen  6  x  i  =  6 
Oxygen  i  x  16  =  16 

Molecular  weight      =  46 

In  alcohol,  therefore,  there  are  12  parts  by  weight  of  carbon, 
6  parts  of  hydrogen,  and  16  of  oxygen.  This  reduced  to 
the  decimal  system  becomes 

Carbon  ^  =  52-17% 
Hydrogen  —  =  13-05  % 
Oxygen  =  34.78  % 


100.00  % 

In  case  there  is  water  of  crystallization  in  the  compound, 
the  water  molecules  combined  with  each  molecule  of  the 
compound  must  also  be  taken  into  account. 

Examples,  i.  Calculate  the  percentage  composition  of 
AgCl;  NaN03;  Ca3(PO4)2;  CHC13;  K4Fe(CN)6;  (NH4)2 
SO4. 

2.  Calculate  the  percentage  composition  of  MgSO4, 
7H2O;  HNaNH4P04,4  H2O. 

Thermochemistry.  It  is  necessary  even  for  the  student  of 
elementary  chemistry  to  understand  at  least  the  fundamental 


LAWS    AND    THEORIES    OF    CHEMISTRY  145 

relations  between  heat  and  chemical  change.  All  chemical 
changes  are  accompanied  either  by  the  using  up  or  the 
giving  out  of  heat.*  For  instance,  when  2  grams  of  hydrogen 
unite  with  18  grams  of  oxygen,  68924  heat  units  are  liberated. 
Such  actions  are  called  exothermic.  There  are  also  chemi- 
cal changes  which  absorb  heat.  For  instance,  to  unite 
i  gram  of  hydrogen  with  127  grams  of  iodin  to  form  hydrio- 
dic  acid,  requires  6000  heat  units.  Such  actions  are  called 
endothermic. 

*  It  is  well  to  note  here  the  relation  between  chemical  energy  and 
other  forms  of  energy.  Energy  is  the  power  that  matter  has  of  doing 
mechanical  work,  i.e.,  of  overcoming  resistance.  Energy  may  be  of  two 
kinds,  either  kinetic  or  potential.  Kinetic  energy  is  the  energy  matter 
has  by  virtue  of  its  motion,  while  potential  energy  is  the  energy  it  has 
by  virtue  of  its  position  or  condition.  For  instance,  a  flying  cannon 
ball  can  overcome  resistance  by  reason  of  its  motion.  On  the  other 
hand,  the  same  cannon  ball,  supported  at  a  height  from  the  ground,  has 
in  it,  by  virtue  of  its  position,  the  power  of  acquiring  motion  when  the 
support  is  removed,  and  of  thus  overcoming  resistance.  The  chemical 
energy  possessed  by  matter  is  potential  energy.  Coal,  on  account  of 
the  affinity  of  carbon  for  oxygen,  will  burn,  and  so  give  out  heat.  This 
heat  in  turn  may  be  utilized  to  boil  water,  the  steam  from  which  in  ex- 
panding will  move  the  piston  of  an  engine,  and  thus  do  mechanical 
work.  Also  the  chemical  energy  of  matter  may  be  transformed  into 
electric  energy  as  in  the  galvanic  cell.  This  electrical  energy  may  be 
utilized  to  drive  a  motor,  and  thus  do  mechanical  work. 

It  has  been  shown  that  carbof?  dioxid  neither  burns  nor  supports 
combustion  (see  Exp.  8  b  ).  The  reason  for  this  is  that  the  carbon  has 
already  united  with  all  the  oxygen  that  it  can  hold;  therefore  no  further 
combustion  can  take  place.  Just  as  the  cannon  ball  would  have  to  be 
lifted  again  to  possess  potential  energy,  so  the  carbon  dioxid  would 
have  to  be  decomposed  in  order  to  possess  chemical  energy  again. 
Carbon  dioxid  does  not  support  combustion,  because  the  affinity  between 
carbon  and  oxygen  is  stronger  than  that  between  oxygen  and  most  other 
elements.  However,  if  an  element  having  a  stronger  affinity  for  oxygen, 
as  for  instance  potassium  at  a  high  temperature,  is  brought  in  contact 
with  carbon  dioxid,  then  the  gas  supports  combustion. 


146       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

We  shall  use  as  our  unit  of  heat  the  calorie,  i.e.,  the 
amount  of  heat  necessary  to  raise  one  gram  of  water  one 
degree.  The  number  of  calories  given  out  by  the  burning 
of  one  gram  of  various  substances  has  been  determined  by  a 
number  of  investigators.  The  following  table  gives  a  few 
examples. 

Hydrogen        34180  Sulfur  2220 

Carbon  8080  Zinc  1300 

Phosphorus      5747  Iron  1181 

We  have  learned  that  the  equation 
2  H2  +  02  =  2  H2  0 

is  in  itself  a  statement  of  the  law  of  the  indestructibility  of 
matter.  But  we  know  that,  in  the  burning  of  hydrogen,  heat 
is  evolved.  This  equation,  as  it  stands,  tells  us  nothing  in 
regard  to  this  fact.  It  is  very  easy,  however,  to  supplement 
the  equation  so  that  it  shall  convey  the  full  meaning.  Let 
the  chemical  symbols  in  an  equation  stand  for  the  number 
of  grams  corresponding  to  its  molecular  weight.  Thus  O2 
stands  for  32  grams  of  oxygen,  HC1  stands  for  35.2  grains 
of  hydrochloric  acid.  By  adding  to  the  chemical  equation 
the  number  of  calories  of  heat  evolved  as  the  result  of  the 
action,  we  express  the  meaning  in  full.  In  the  example  of 
burning  hydrogen  we  have 

2  H2  +  02  =  2  H2  0  +  136800 

which  signifies  that  4  grams  of  hydrogen  on  uniting  with  32 
grams  of  oxygen  form  36  grams  of  water,  and  at  the  same 
time  give  out  136800  calories  of  heat.  In  the  case  of  en- 
dothermic  reactions,  the  notation  is  the  same,  except  that  the 
—  sign  instead  of  the  +  sign  is  used.  In  addition  to  this, 
heavy  type  is  used  to  represent  solids,  ordinary  type  liquids, 
and  italics,  gases.  Thus, 

c  +  a,  =  c  a,  +  97000 


LAWS    AND    THEORIES    OF    CHEMISTRY 

signifies  that  12  grams  of  solid  carbon  burned  in  32  grams 
of  gaseous  oxygen  gives  44  grams  of  gaseous  carbon  dioxid 
together  with  97000  calories  of  heat. 

In  a  great  many  cases  where  the  heat  liberated  cannot  be 
found  by  experiment,  it  can  be  found  by  calculation. 

Example.  Required  to  find  the  heat  liberated  when  car- 
bon is  oxidized  to  carbon  monoxid.  It  is  found  that  when 
12  grams  of  carbon  is  oxidized  to  carbon  dioxid,  97000* 
calories  are  liberated,  and  that  when  28  grams  of  carbon 
monoxid  is  oxidized  to  carbon  dioxid,  136000  |  calories  are 
liberated.  We  then  have 

C  +  6>2  =  C  O2  +  97000 
2  C  O  +  O-2  —  2  C  O.2  +  136000 

By  dividing  the  second  equation  by  2,  we  obtain  68000  cal. 
as  the  amount  of  heat  liberated  when  44  grams  of  carbon 
dioxid  are  formed.  We  divide  by  2  in  order  to  obtain  the 
same  weight  of  carbon  dioxid  as  in  the  first  equation.  Sub- 
tracting 68000  from  97000,  we  evidently  obtain  the  number 
of  calories  given  out  when  carbon  is  oxidized  to  the  monoxid, 
i.e.,  29000  cal.  We  therefore  have  the  equation 

C  +  O  =  CO  +  29000 

EXPERIMENT    21 
Heat  of  Chemical  Action 

Dilute  50  cc.  of  concentrated  sulfuric  acid  with  250  cc. 
of  water,  and  allow  it  to  cool.  Weigh  a  beaker  of  thin  glass 
large  enough  to  hold  the  acid,  and  place  it  in  a  larger  beaker. 
Pack  wool  around  and  under  the  inner  beaker  in  such  a  way 
that  the  inner  beaker  can  be  easily  removed.  Let  this  inner 
beaker  be  used  as  a  calorimeter.  Pour  the  acid  into  the 
*  More  correctly  96960.  t  More  correctly  135920. 


148        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

calorimeter.  Clean  a  piece  of  sheet  zinc  about  2\  inches 
wide  by  5  long.  Take  the  temperature  of  the  acid,  then 
plunge  the  zinc  into  it.  When  the  temperature  has  risen 
four  or  five  degrees,  remove  the  zinc,  stir,  and  take  the  tem- 
perature carefully.  Wash  the  zinc  clean,  dry  it,  and  weigh 
again. 

It  will  be  necessary  to  know  the  specific  heat  of  the  solu- 
tion after  the  zinc  has  been  acted  upon.  To  find  this,  per- 
form an  experiment  exactly  like  Experiment  15,  Part  II.  with 
the  exception  of  using  this  calorimeter  and  liquid  just  as 
they  are,  instead  of  the  lemonade  shaker  and  water.  Calcu- 
late the  specific  heat  of  the  liquid,  knowing  the  specific  heat 
of  lead  to  be  .031. 

Arrange  your  calculations  as  follows  :  — 

To  find  the  specific  heat  of  the  liquid. 

Wt.  of  liquid  =  1  — 
Wt.  of  calorimeter  =  c  = 
Wt.  of  shot  =  s  = 
Specific  heat  of  shot  =  .031 
Specific  heat  of  glass  =  .2 
Temperature  liquid  =  t  = 
Temperature  liquid  and  shot  =  t'  = 
Temperature  of  shot  =  t"  = 

Let  x  =  specific  heat  of  the  liquid. 

.031  s  (t"  -  t')  =  .2  c  (f  -  t)  +  xl  (f  - 1) 
Solve  for  x. 

To  find  the  amount  of  heat  liberated  when  one  gram  of 
zinc  is  dissolved  in  sulfuric  acid. 

Wt.  of  zinc  dissolved  —  z  = 

Wt.  of  water  and  acid  =  1  = 

Wt.  of  calorimeter  =  c  = 

Specific  heat  of  glass  =  .2 

Specific  heat  of  liquid  =  s  (found  in  foregoing  calculation) 

Temperature  before  action  =  t  = 

Temperature  after  action  ='?  = 


LAWS    AND    THEORIES    OF    CHEMISTRY  149 

Let  x  =  the  number  of  calories  liberated  when  one  grm. 
of  zinc  is  dissolved. 

The  amount  of  heat  evolved  by  the  action  of  the  acid  on 
the  zinc  equals  the  amount  of  heat  gained  by  the  solution 
and  the  calorimeter. 

X  Z  =  Sl  (f  -  t)  +  .2  C  (f  -  t) 

Solve  for  x  and  obtain  the  number  of  calories  liberated 
when  one  gram  of  zinc  is  dissolved.  Find  65^  in  con- 
formity with  principle  laid  down  on  page  146. 

Write  the  reaction  for  the  chemical  change  together  with 
the  heat  evolved. 

EXPERIMENT   22 
Heat  of  Neutralization 

Weigh  out  about  25  grams  of  c.p.  concentrated  sulfuric 
acid.  Dilute  this  with  200  cc.  of  water,  and  allow  it  to  cool. 
Calculate  (see  page  142)  the  number  of  grams  of  sodium 
hydroxid  necessary  to  neutralize  the  acid,  take  about  one- 
sixth  more  than  this  amount,  dissolve  it  in  200  cc.  of  water, 
and  allow  the  solution  to  cool.  Let  the  two  vessels  con- 
taining respectively  the  acid  and  alkali  stand  side  by  side 
until  they  are  of  the  same  temperature.  Then  pour  them 
together  in  the  glass  calorimeter  used  in  Exp.  21,  and  note 
the  rise  in  temperature. 

Arrange  your  calculations  as  follows  :  — 

Wt.  of  sulfuric  acid  =  s  = 
Wt.  of  sodium  hydroxid  =  h  = 
Wt.  of  calorimeter  =  c  = 
Specific  heat  of  glass  =  .2 
Temperature  before  mixing  =  t  = 
Temperature  after  mixing  =  t'  = 
Wt.  of  water  =  w  = 
Specific  heat  of  solution  =  a  = 


I5O       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Find  a,  the  specific  heat  of  the  sodium  sulfate  solution,  in 
the  same  way  as  you  did  that  of  zinc  sulfate  in  Experiment 
2 1 .  The  heat  gained  by  the  solution  and  calorimeter  equals 
the  heat  evolved  by  neutralization. 

Heat  gained  by  solution  =  (s  +  w  +  h)  (f  —  t)  a 
Heat  gained  by  calorimeter  =  .2  c  (f  —  t) 
Heat  of  neutralization  =  sx 
Therefore 

(s  +  w  +  h)  (t  '-  t)  a  +  .2  c  (f  -  t)  =  sx 

x  = 
98  x  = 

We  find  98^  in  conformity  with  the  principle  laid  down  on 
page  146. 

Write  the  reaction  for  the  chemical  change  together  with 
the  heat  evolved. 

EXPERIMENT    23 
Heat  of  Solution  and  of  Hydration 

a.  Place  in  a  weighed  calorimeter  about  350  grams  of 
water  weighed  accurately  to  one  gram.  Powder  exactly  40 
grams  of  anhydrous  magnesium  sulfate,  place  it  in  a  beaker, 
and  cover  it  with  a  watch  glass.  Allow  the  calorimeter, 
containing  the  water,  and  the  beaker  to  stand  side  by  side 
until  they  are  of  the  same  temperature.  Then  pour  the 
powdered  salt  into  the  water,  and  stir  the  solution  with  a 
thermometer,  taking  the  temperature  when  all  is  dissc'ved. 

Arrange  your  calculations  as  follows  :  — 

Wt.  of  magnesium  sulfate  —  m  = 
Wt.  of  calorimeter  =  c  = 
Wt.  of  water  =  w  = 
Temperature  before  mixing  =  t  = 
Temperature  after  mixing  =  t'  = 
Specific  heat  of  solution  =  a  = 
Specific  heat  of  glass  =  .2 


LAWS    AND    THEORIES    OF    CHEMISTRY  151 

Find  a  in  the  same  way  as  in  Experiments  2 1  and  2  2 . 

The  heat  here  given  out  is  made  up  of  two  different  heats  ; 
first,  the  heat  given  out  by  the  anhydrous  salt  taking  on 
water  of  crystallization ;  second,  the  heat  used  up  by  the 
salt  dissolving  in  water. 

We  then  have  the  equation  :  — 

Heat  gained  by  the  solution  and  calorimeter  =  heat  of 
hydration  and  of  solution. 

Heat  gained  by  the  solution  =  (w  -f  m)  (f  —  t)  a 
Heat  gained  by  the  calorimeter  =  .2  c  (t'  —  t) 
Heat  of  hydration  and  of  solution  =  mx 

Therefore 

(w  +  m)  (f  -  t)  a  +  .2  c  (f  - 1)  =  mx 

x  = 

I2O  X  = 

The  number  120  is  taken  for  the  same  reason  that  98  was 
taken  in  the  last  experiment.  (See  page  146.) 

You  should  obtain  as  a  result  somewhere  near  the  number 
20280. 

b.  Now  repeat  the  operation,  using  exactly  82  grams  of 
the  salt  with  its  water  of  crystallization  (MgSO4,7H2O). 
In  this  case,,  calculate  the  heat  lost  in  dissolving  246  grams 
of  magnesium  sulfate.  You  should  obtain  a  number  some- 
where near  3800. 

We  see  then  that  the  heat  of  hydration  (a  positive  heat) 
must  be  3800  greater  than  the  heat  of  solution  (a  negative 
heat),  i.e.,  24080. 

Dissociation.  We  have  found  (Exp.  42,  Part  I.)  that  when 
ammonium  chlorid  is  heated  it  passes  directly  from  the  solid 
to  the  gaseous  state.  If  the  density  of  ammonium  chlorid  is 
found  in  the  state  of  vapor,  the  number  obtained  is  13.34. 
Now,  according  to  the  hypothesis  of  Avogadro,  the  number 


I$2         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

should  be  26.69,  smce  the  molecular  weight  is  53.38.  The 
densities  of  hydrochloric  acid  gas  and  ammonia  are,  how- 
ever, 18.18  and  8.5  respectively;  which,  if  added,  give  26.68, 
showing  that  undoubtedly,  when  ammonium  chlorid  is  vapor- 
ized, it  is  not  made  up  of  molecules  of  ammonium  chlorid, 
but  of  a  mixture  of  molecules  of  hydrogen  chlorid  and 
ammonia.  When  allowed  to  cool,  these  molecules  reunite 
and  form  molecules  of  the  original  compound.  This  pheno- 
menon is  called  dissociation. 


EXPERIMENT    24 
Dissociation 

Procure  a  glass  tube  about  25  cm.  long  and  2  cm.  bore, 
and  close  the  ends   by  means  of  two  closely  fitting  corks, 

through  both  of  which  passes 
the  stem  of  a  clay  tobacco 
pipe.  Place  in  the  center  of 
the  tube  a  piece  of  crystal- 
lized ammonium  chlorid,  and 
at  the  ends,  next  to  the  corks, 
place  pieces  of  moist  blue 
litmus  paper,  flat  against 
the  glass.  Connect  one 
end  of  the  pipe  with  a  pair 
of  bellows.  Heat  the  am- 
monium chlorid  with  the 
flame  of  a  Bunsen  burner, 
at  the  same  time  gently 
forcing  air  through  the  pipe 
stem.  The  two  gases  into 
which  the  ammonium  chlorid  has  been  dissociated  will  pass 
through  the  porous  pipe  stem  in  different  quantities;  and, 


LAWS  AND  THEORIES  OF  CHEMISTRY       153 

by  holding  a  piece  of  moist  red  litmus  paper  at  the  opening 
of  the  pipe,  the  presence  of  ammonia  gas  will  be  shown, 
while  the  paper  inside  the  tube  will  show  the  presence  of 
free  hydrochloric  acid  gas. 

Dissociation  in  Solutions.  Dissociation  takes  place  not  only 
in  gaseous  compounds  but  also  in  solutions.  In  these  cases 
it  is  probable,  and  it  is  believed,  that  the  molecules  are  con- 
tinually breaking  up  into  atoms  or  groups  of  atoms,  and 
then  reuniting  again.  These  atoms  or  groups  of  atoms  are 
called  ions.  In  the  case  of  a  solution  of  hydrochloric  acid, 
the  ions  would  be  H  and  Cl,  and  in  the  case  of  a  solution  of 
copper  sulfate  they  would  be  Cu  and  SO4.  The  ions  are 
believed  to  be  present  in  greater  numbers  in  dilute  solutions 
than  in  strong  ones.  The  theory  is  that  the  difference  be- 
tween ions  and  simple  atoms  is,  that  the  ions  are  charged 
electrically ;  thus,  in  a  solution  of  copper  sulfate,  the  Cu 
ions  are  charged  positively,  and  the  SO4  ions  are  charged 
negatively. 

EXPERIMENT    25 
Dissociation  in  Liquids 

Prepare  four  strong  aqueous  solutions  (10  cc.  each)  as  fol- 
lows :  —  one  of  copper  nitrate,  one  of  copper  chlorid,  one  of 
sodium  chlorid,  and  one  of  sodium  nitrate.  In  a  test  tube, 
mix  the  copper  nitrate  and  the  sodium  chlorid  solutions.  In 
another,  mix  the  sodium  nitrate  and  copper  chlorid.  Note 
that  the  two  mixtures  are  the  same.  Evidently  each  mixture 
contains  copper  nitrate,  copper  chlorid,  sodium  nitrate,  and 
sodium  chlorid.  Besides  these  compounds  there  must  be 
present  the  ions  of  Cu,  Na,  NO3  and  Cl.  To  destroy  the 
equilibrium  between  these,  add  to  one  of  the  mixtures  a  little 


i$4      AN  ELEMENTARY  EXPERIMENTAL  CMEMISTRV 

powdered  sodium  chlorid  and  shake.  The  additional  green 
color  shows  the  presence  of  more  copper  chlorid,  the  forma- 
tion of  which  has  of  course  necessitated  a  rearrangement  of 
the  compounds  and  ions. 

Organic  Chemistry.  It  was  formerly  thought  that  those 
compounds  formed  by  the  chemical  elements  in  living  bodies 
were  not  bound  by  the  same  laws  as  those  of  the  inorganic 
world.  The  principle  of  life  was  supposed  in  some  mys- 
terious way  to  govern  them,  and  it  was  thought  that  they 
could  not  be  prepared  artificially.  On  this  account,  that 
branch  of  chemistry  which  deals  with  such  compounds  was 
named  organic  chemistry.  When  Wohler,  a  German  chemist, 
succeeded  in  making  artificially  the  organic  compound  urea, 
this  theory  was  overthrown.  After  many  other  compounds 
were  thus  made,  it  became  no  longer  tenable,  and  we  now 
class  all  chemical  compounds  as  dependent  upon  the  same 
laws.  It  happens  that  the  element  carbon  is  the  most  fre- 
quently present  in  the  so-called  organic  compounds ;  so  it 
were  better  to  name  this  branch  of  the  science  the  chemistry 
of  the  compounds  of  carbon.  However,  the  name  organic 
chemistry  has  clung  to  it,  and  probably  always  will. 

EXPERIMENT    26 
Burning  of  Organic  Matter;  Dry  Distillation 

a.  Burn  a  number  of  organic  substances  such  as  wood, 
alcohol,  a  candle,  kerosene,  etc.,  under  a  cold  bell-jar,   and 
notice  the  formation  of  water.     On  the  end  of  a  glass  rod 
hold  a  drop  of  lime-water  in  the  jar,  and  prove  the  presence 
of  carbon  dioxid. 

b.  To  a  hard-glass  tube  about  8  mm.  in  diameter  con- 
taining a  piece  of  wood,  attach  by  means  of  a  one-holed  cork 


LAWS    AND    THEORIES    OF    CHEMISTRY  155 

a  glass  exit  tube.  Heat  the  wood,  and  ignite  the  gas  that 
escapes.  What  remains  in  the  tube  ?  Is  the  combustion  of 
the  wood  partial  or  complete  ? 

Alcohol.  Sugar  is  a  compound  made  up  of  carbon,  hydro- 
gen, and  oxygen  (CgH^Oe).  Whenever  a  juice  containing 
sugar  is  left  in  the  open  air,  it  decomposes,  giving  off 
carbon  dioxid  gas,  and  forming  a  new  compound  called 
alcohol  C2H6O. 

C6  Hi2  06  =  2  C2  He  0  +  2  C  02 

This  action,  called  fermentation,  is  caused  by  a  small 
organized  body  (in  this  case  vegetable)  called  a  ferment. 

EXPERIMENT    27 
Alcohol 

In  a  500  cc.  flask,  dissolve  40  grms.  of  grape  sugar  in  250 
cc.  of  water.     Add  to  this  a  little  brewer's  yeast,  after  con- 
necting the   flask  with  a  wash  bottle 
containing   lime   water.       Notice  the 
evolution  of  carbon  dioxid,  as  proved 
by  the  milky  color  of  the  lime  water. 
After  the   apparatus  has   stood  long 
enough  for  the   action  to  cease,   re- 
move the  flask  and  place  it  in  a  water 
bath.     Connect  it  with  condenser,  and 
allow  the  alcohol  to  distill  over.    The 
condenser  may  be  made  as  follows. 
Fit  each  end  of  a  glass  tube  (about  60  cm.  long  and  2.5 
or  3  cm.  bore)  with  a  two-holed  rubber  stopper.     Let  a  glass 
tube  about  75  cm.  long  and  5   mm.  bore  extend  through  the 
tube,  and  through  both  stoppers.     In  the  other  holes  of  the 


156       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

stoppers,  fit  pieces  of  glass  tubing  about  10  cm.  long  and 
bent  at  right  angles.  Fasten  the  apparatus  to  a  stand,  and 
incline  it  at  a  slight  angle.  Allow  cold  water  by  means  of 
rubber  tubing  to  enter  the  lower  end  and  escape  from  the 
higher. 


EXPERIMENT    28 
Saponification 

In  a  porcelain  dish,  boil  for  an  hour  about  one-eighth  of  a 
pound  of  lard  together  with  a  solution  of  sodium  hydroxid 
(20  grms.  to  125  cc.  of  water).  While  the  mixture  is  cooling, 
add  a  strong  solution  of  salt.  The  substance  that  solidifies 
on  the  surface  of  the  liquid  is  soap. 

Remark.  Soaps  are  the  alkali  salts  of  certain  fatty  organic 
acids  such  as  stearic,  C18H36O2,  and  palmitic  C16H32O2.  The 
calcium  and  magnesium  salts  of  these  acids  are  insoluble 
in  water;  hence,  when  soap  is  used  with  hard  water  (see 
Exp.  35,  Part  I.),  these  compounds  appear  on  the  surface 
of  the  water  as  a  scum. 


PART  III. 

HISTORY,  OCCURRENCE  AND 

INDUSTRIAL  APPLICATIONS  OF  THE 

PRINCIPAL  ELEMENTS  AND  COMPOUNDS 


PART   III. 

HISTORY,  OCCURRENCE  AND   INDUSTRIAL 

APPLICATIONS    OF    THE    PRINCIPAL    ELEMENTS 

AND   COMPOUNDS 


OXYGEN 

History.  Until  1774,  the  air  was  believed  to  be  a  simple 
substance.  In  that  year,  the  investigations  of  Priestley, 
Rutherford,  and  Scheele  proved  that  it  was  a  mixture  of  two 
different  gases.  By  heating  mercuric  oxid  (see  Exp.  5, 
Part  I.),  Priestley  proved  that  this  substance  was  composed 
of  a  gas  and  metallic  mercury.  The  gas  thus  obtained  was 
shown  to  be  the  same  as  one  of  the  constituents  of  the  air. 
In  1805,  Gay  Lussac  proved  that  water  was  composed  of  two 
volumes  of  hydrogen  and  one  of  oxygen.  (See  Exp.  6  e, 
Part  I.)  The  name  oxygen  (6£v's,  sour,  yewaw,  I  produce) 
was  given  it  by  Lavoisier. 

Occurrence.  Oxygen  is  the  most  abundant  element  in 
nature.  It  constitutes  23  per  cent  of  the  atmosphere,  88.88 
per  cent  of  water,  and  from  44  to  48  per  cent  of  the  crust  of 
the  earth.  The  ores  of  almost  all  of  the  metals  occur  in  the 
earth  as  oxids,  or  as  other  compounds  containing  oxygen. 

Industrial  Applications.  Oxygen  is  used  extensively  for 
medical  purposes,  and  in  connection  with  hydrogen  in  the 
oxy-hydrogen  blowpipe.  It  is  usually  prepared  in  large 
quantities  by  heating  potassium  chlorate  (KC1O8),  or  by 
obtaining  it  from  the  atmosphere.  In  the  latter  method, 


I6O        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

advantage  is  taken  of  the  fact  that  at  a  dull  red  heat  barium 
oxid  (BaO)  takes  on  oxygen,  becoming  barium  peroxid 
(BaO2),  and  that  at  a  still  higher  temperature  the  oxygen 
thus  absorbed  is  given  off,  leaving  the  original  oxid.  Theoret- 
ically this  process  could  be  continued  indefinitely.  In  prac- 
tice, however,  the  barium  oxid  is  not  used  indefinitely,  since 
it  becomes  gradually  less  efficient  in  its  action.  The  changes 
are 

2  Ba  0  +  02  =  2  Ba  02 

2  Ba  02  =  2  Ba  0  +  02 

It  has  been  found  that  this  reaction  can  be  accomplished  at 
a  constant  temperature  by  changing  the  pressure. 

HYDROGEN 

History.  In  the  sixteenth  century.  Paracelsus  obtained  an 
inflammable  gas  by  treating  metals  with  certain  acids.  To 
this  gas,  Cavendish  gave  the  name  "  Inflammable  Air  "  in 
1766.  Later  he  proved  that,  when  this  gas  was  united  with 
oxygen,  water  was  formed.  Lavoisier  confirmed  this,  and 
gave  to  the  gas  the  name  hydrogen  ({SStop,  water,  yevraw,  I 
produce). 

Occurrence.  Hydrogen  occurs  in  the  free  state  in  the 
atmosphere  of  the  sun,  in  small  quantities  mixed  with  other 
gases  in  volcanic  eruptions,  and  sometimes  in  oil  wells.  It 
is  also  found  occluded  in  meteoric  iron  and  certain  iron  ores. 
In  chemical  combination,  it  is  most  widely  distributed  as 
water.  It  also  occurs  in  combination  with  a  number  of  the 
non-metals,  and  as  a  part  of  almost  all  organic  compounds. 

Industrial  Applications.  From  its  great  lightness,  hydrogen 
is  valuable  for  filling  balloons.  It  has  often  been  made  for 
military  balloons  by  the  steam  and  red-hot  iron  process. 
(See  Exp.  15  b,  Part  I.) 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          l6l 

The  great  amount  of  heat  given  out  by  the  oxidation  of 
hydrogen  has  made  it  valuable  in  melting  refractory  metals 
such  as  platinum.  This  is  done  by  means  of  a  very  simple 
piece  of  apparatus  called  the  oxy-hydrogen  blowpipe.  It  is 


simply  a  tube  within  a  tube,  the  tip 
being  made  of  platinum.  The  inner 
tube  is  connected  with  a  gas  holder 
containing  hydrogen,  while  the  other  is  connected  with  a 
similar  one  containing  oxygen.  The  hydrogen  is  first  turned 
on  and  ignited.  The  oxygen  is  then  turned  on  until  the 
flame  burns  quietly. 

WATER 

Water  is  one  of  the  most  abundant  and  most  widely 
distributed  compounds.  From  the  fact  that  it  is  a  solvent 
for  so  many  substances,  it  is  never  found  pure.  When 
pure,  it  is  tasteless  and  colorless.  In  large  quantities,  it 
often  has  a  greenish  or  a  bluish  color.  The  great  reservoir 
of  course  is  the  sea,  from  which  the  water  that  is  pre- 
cipitated upon  the  land  as  rain,  hail,  or  snow,  originally 
evaporated.  Sea  water  contains  in  solution  about  3^  per 
cent  of  solid  matter,  most  of  which  is  sodium  chlorid.  The 
water  of  lakes  and  rivers  contains  various  substances  depend- 
ing upon  the  locality.  The  most  common  substances  are  cal- 
cium carbonate  and  calcium  sulfate.  (See  Exp.  31,  Part  I.) 

Water  is  so  widely  used  that  it  is  hardly  necessary  to 
enumerate  many  of  the  ways  in  which  it  is  of  value. 
The  principal  uses  of  water,  however,  are  for  drinking, 


1 62       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

washing,  and  culinary  purposes  in  the  household,  and  for 
boiling  and  solvent  purposes  in  the  arts.  Good  drinking 
water  should  be  as  free  as  possible  from  sewage  contamina- 
tion and  decaying  organic  matter.  These  contaminations 
are  not  only  injurious  in  themselves,  but  they  render  the 
water  more  habitable  for  the  germs  of  typhoid  fever,  cholera, 
and  other  malignant  diseases.  The  disease  germs  are  almost 
always  carried  into  water  by  means  of  sewage,  hence  water 
containing  the  slightest  trace  of  sewage  should  be  especially 
avoided.  Water  that  contains  organic  impurities  is  usually 
yellowish  in  color,  and  has  a  disagreeable  odor. 

The  water  supply  of  large  cities  is  freed  from  impurities 
by  a  process  of  filtration  through  several  feet  of  sand  and 
broken  stone  extending  over  a  large  area.  At  least  two  such 
filters  are  used,  and  one  is  kept  empty  while  the  other  is  in 
use.  The  disease  germs  that  are  removed  by  one  filter  are 
destroyed  by  oxidation,  when  the  water  is  diverted  into  the 
other. 

For  boiling  and  solvent  purposes  in  the  arts,  soft  water  is 
preferable.  For  washing  purposes,  hard  water  may  be  made 
soft  by  the  addition  of  sodium  carbonate. 

HYDROGEN    DIOXID,  H2O2 

This  compound  was  first  discovered  by  Thenard  in  1818. 
He  obtained  it  by  treating  barium  peroxid  with  dilute  hydro- 
chloric acid. 

Ba  02  +  2H  Cl  =  H2  02  +  Ba  C12 

It  is  formed  in  minute  traces  in  the  atmosphere,  and  is 
sometimes  produced  simultaneously  with  the  preparation  of 
ozone.  It  is  an  oily,  colorless  liquid,  having  a  bitter  taste. 
It  is  usually  used  in  dilute  solutions. 

Industrial  Applications.     Because  of  the  fact  that  hydro- 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          163 

gen  peroxid  gives  up  part  of  its  oxygen  readily,  it  is  very 
valuable  as  an  oxidizing  agent.  It  has  been  extensively 
used  for  bleaching  hair,  giving  to  dark  hair  the  well-known 
light  flaxen  tint.  It  is  also  used  for  restoring  the  colors  to 
old  paintings.  White  paint  is  composed  largely  of  lead  car- 
bonate, which  gradually  darkens  in  time.  This  occurs  on 
account  of  the  action  of  sulfur,  which  forms  a  black  lead  sul- 
fid.  If  hydrogen  peroxid  is  applied  to  such  a  discolored 
painting,  the  lead  sulfid  is  converted  into  lead  sulfate,  thus 
restoring  in  a  great  measure  the  original  color.  Hydrogen 
peroxid  is  also  used  extensively  in  medicine  and  as  a 
disinfectant. 

THE   HALOGENS 

The  elements  chlorin,  bromin,  iodin,  and  fluorin  may  be 
grouped  together,  since  their  properties  are  in  many  ways 
similar.  This  group  is  called  the  halogen  group  (oAos,  salt, 
yewda),  I  produce). 

CHLORIN 

History.  Chlorin  gas  was  first  obtained  and  studied  by 
Scheele  in  1774.  He  obtained  it  from  hydrochloric  acid 
and  manganese  dioxid.  Its  elementary  character  was  proved 
by  Davy  in  1810,  and  he  named  it  chlorin  from  ^Xwpos, 
meaning  greenish  yellow. 

Occurrence.  On  account  of  its  strong  affinity  for  other 
elements,  chlorin  is  not  found  in  the  free  state.  It  occurs  in 
combination  chiefly  with  the  alkali  metals  in  sea  water,  and 
as  rock  salt  in  various  localities  (chiefly  at  Stassfurt,  Ger- 
many). At  Syracuse,  New  York,  salt  occurs  as  a  brine  at  a 
depth  of  from  200  to  400  feet  below  the  surface  of  the  earth. 

Industrial  Applications.  From  its  strong  affinity  for  hydro- 
gen, chlorin  is  used  in  the  arts  in  the  process  of  bleaching 


164       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

cloth.  In  bleaching  cloth,  it  is  necessary  for  the  cloth  to  be 
moist.  The  change  is  explained  by  the  fact  that  the  chlorin 
unites  with  the  hydrogen  of  the  water,  leaving  nascent 
oxygen,  which  in  turn  oxidizes  the  vegetable  coloring  matter 
into  colorless  compounds.  Black  colors  from  carbon,  as  for 
instance  printer's  ink,  cannot  be  bleached.  Ordinary  writing 
ink,  which  is  a  compound  of  organic  acids  and  iron,  is  readily 
decolorized. 

Large  quantities  of  chlorin  are  also  used  for  disinfectant 
purposes. 

Manufacture.  It  will  be  remembered  that  chlorin  is  made 
in  the  laboratory  (see  Exp.  27  b)  by  treating  manganese 
dioxid  with  hydrochloric  acid,  thus  :  — 

Mn  02  +  4  H  Cl  =  Mn  Cls  +  2H2  0  +  C12 

Since  manganese  dioxid  is  the  costly  material  used,  it  is 
necessary  in  the  manufacture  of  chlorin  on  a  large  scale  to 
save  the  manganese  chlorid  formed,  and  in  some  way  change 
it  back  to  the  oxid.  In  the  Weldon  process  this  is  done. 
The  chlorid  is  treated  with  calcium  hydrate.  The  mixture 
is  then  heated,  and  a  current  of  air  is  blown  through.  A 
complicated  reaction  that  need  not  be  entered  into  here 
ensues,  and  the  manganese  is  oxidized  to  peroxid.  Thus 
the  original  manganese  can  be  used  over  again,  making 
chlorin  »a  cheap  commercial  product. 

Other  processes  for  the  manufacture  of  chlorin  are  also  in 
use. 

HYDROCHLORIC    ACID 

History  and  Occurrence.  Hydrochloric  acid  was  known  to 
Arabian  alchemists  in  aqua  regia.  Basil  Valentine  in  the 
fifteenth  century  wrote  of  an  acid  which  he  called  "  spirit  of 
salt,"  obtained  from  oil  of  vitriol  and  common  salt.  Priestley 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         165 

obtained  the  gas  by  collecting  it  over  mercury  instead  of 
water,  and  called  it  "marine  add  air"  In  1810,  Davy 
showed  that  this  gas  was  composed  of  hydrogen  and  chlorin, 
and  that  it  was  not  a  compound  of  oxygen  as  had  formerly 
been  supposed. 

Hydrochloric  acid  occurs  in  the  gases  which  issue  from 
some  volcanoes,  especially  Vesuvius.  It  is  also  found  in 
some  South  American  rivers  whose  sources  are  in  the  vol- 
canic districts  of  the  Andes. 

Industrial  Applications  and  Manufacture.  Hydrochloric 
acid  is  used  in  the  manufacture  of  chlorin,  ammonium  chlorid, 
and  tin  chlorid,  the  last  being  extensively  used  by«dyers. 

In  the  manufacture  of  sodium  sulfate,  large  quantities  of 
hydrochloric  acid  are  obtained  as  a  by-product  (see  Exp.  23, 
Part  I.).  The  acid  fumes  pass  through  a  flue  to  brick 
chambers  filled  with  coke  or  broken  brick  through  which 
water  is  passing.  The  water  absorbs  the  gas,  and  this  solu- 
tion, when  collected,  is  the  hydrochloric  acid  of  commerce. 

OXIDS    AND    OXY-ACIDS    OF   CHLORIN 

There  are  three  oxids  of  chlorin. 

Chlorin  monoxid,  C12  0 
Chlorin  trioxid,  C12  03 
Chlorin  peroxid,  Cl  02 

They  are  all  unstable  compounds. 
The  oxy-acids  of  chlorin  are 

Hypochlorous  acid,  H  Cl  0 
Chlorous  acid,  H  Cl  02 
Chloric  acid,  H  Cl  03 
Perchloric  acid,  H  Cl  02 

These  only  exist  in  aqueous  solutions.  Hypochlorous  and 
chloric  acids  form  salts  that  are  important  commercial 
products. 


1 66         AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


BLEACHING   POWDER 

If  chlorin  is  allowed  to  pass  into  chambers  containing 
quantities  of  freshly  slaked  lime,  it  is  absorbed  by  the  lime, 
and  a  compound  commonly  called  "  chlorid  of  lime "  *  is 
formed.  This  compound  has  strong  bleaching  properties 
because  it  slowly  gives  off  chlorin.  Considerable  time  has 
been  given  to  the  study  of  the  composition  of  this  substance, 
but  at  present  no  entirely  satisfactory  explanation  has  been 
given.  It  was  first  thought  to  be  a  salt  of  hypochlorous 
acid,  but  it  does  not  show  the  proper  percentage  of  chlorin. 
The  best  authority  on  the  subject  affirms  that  its  composi- 
tion is  CaOCl2,  and  calls  it  chloro-hypochlorite.  It  is  also 
called  calcium  oxychlorid. 

"  Chlorid  of  lime  "  is  used  in  enormous  quantities  as  a 
disinfectant,  and  for  bleaching  purposes.  In  bleaching  calico 
and  paper  pulp,  a  two  per  cent  solution  of  bleaching  powder 
is  used.  The  soaked  product  is  then  placed  in  a  dilute 
solution  of  sulfuric  acid,  which  liberates  the  chlorin  more 
freely.  When  thoroughly  bleached,  the  product  is  treated 
with  sodium  sulfite  (called  anti-chlor),  which  removes  all 
traces  of-  unused  chlorin.  This  is  done  because  otherwise 
the  chlorin  would  slowly  attack  the  fiber  of  the  cloth  or  paper. 

POTASSIUM    CHLORATE 

The   potassium   salt  of  chloric   acid  is  a  very  important 
product.    It   can  be   made  by  treating  warm  potassium  hy- 
droxid  with  chlorin  gas.     The  reaction  is  as  follows  :  — 
3  C12  +6KOH  =  KC103  +  5KC1  +  3H20 

By  this  method,  five   of  the   six   molecules   of  potassium 

*  If  bleaching  powder  is  treated  with  an  acid,  the  chlorin  is  given  off 
again. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         l6/ 

hydroxid  are  converted  into  potassium  chlorid ;  so  it  is  better 
first  to  form  the  corresponding  calcium  salts,  and  then  add 
to  the  solution  enough  potassium  chlorid  to  change  the  cal- 
cium chlorate  to  potassium  chlorate.  The  use  of  potassium- 
hydroxid,  which  is  the  expensive  compound  in  the  first 
method,  is  avoided  in  the  second. 
The  reactions  are 

6  Ca  (0  H)2  +  6  C12  =  Ca  (Cl  03)2  +  5  Ca  C12  +  6  Ho  0 
and 

Ca  (Cl  <V2  +  2  K  Cl  =  2  K  Cl  Os  +  Ca  Cl, 

BROMIN 

History.  Balard  discovered  the  element  bromin  in  1826. 
He  prepared  it  from  bittern,  the  liquid  remaining  after  sodium 
chlorid  has  been  crystallized  out  of  concentrated  sea  water. 
He  named  it  bromin  from  /8poi/xos,  a  bad  odor. 

Occurrence.  Bromin,  as  has  been  stated,  occurs  in  sea 
wrater.  It  also  occurs  in  the  waters  of  many  mineral  springs. 
In  these  cases  it  is  in  the  form  of  bromids  of  sodium,  potas- 
sium, magnesium,  or  calcium.  It  also  occurs  in  ores  in  com- 
bination with  silver,  as  silver  bromid. 

Industrial  Applicatiojt.  The  chief  use  of  bromin  is  in  the 
manufacture  of  bromids  and  as  an  oxidizing  agent.  Its  com- 
pounds are  also  used  in  photography  and  medicine. 

HYDROBROMIC    ACID 

Preparation.  In  Exp.  280,  Part  I.,  the  student  made 
hydrobiomic  acid  by  the  direct  union  of  hydrogen  and 
bromin.  In  e  of  the  same  experiment,  it  was  found  that  the 
acid  could  not  be  prepared  in  a  manner  similar  to  that  of 
hydrochloric  acid.  The  best  way  to  prepare  it  is  to  make 
use  of  the  action  of  phosphorus  bromid  upon  water.  A 


1 68      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

mixture  of  one  part  of  amorphous  phosphorus  and  two  parts 
of  water  is  placed  in  a  flask  fitted  with  a  delivery  tube  and  a 
funnel  tube  with  a  stop-cock.  The  flask  is  connected  with 
a  U  tube  containing  a  mixture  of  broken  glass  and  amor- 
phous phosphorus.  Bromin  is  allowed  to  enter  the  flask, 
drop  by  drop,  and  hydrobromic  acid  is  evolved.  The  phos- 
phorus in  the  U  tube  takes  up  any  bromin  vapors  that  may 
accompany  the  acid  gas. 

P  Br3  +  3  H2  0  =  3  H  Br  +  H3  P  03 


IODIN 

History.  Courtois,  in  1812,  discovered  iodin  in  the  solu- 
tions of  the  sodium  salts  obtained  from  kelp,  the  ashes  of 
seaweed.  Its  name  is  derived  from  teto8rjs,  meaning  violet- 
colored,  on  account  of  the  violet  color  of  its  vapor. 

Occurrence.      Iodin, 

":^^ 


^ke  chlorin  and  bromin, 
does  not  occur  uncom- 
bined  with  other  ele- 
ments. It  is  found  com- 
bined with  the  metals 
in  the  form  of  iodids, 
both  in  the  animal  and 
the  vegetable  kingdoms. 
The  chief  source  is  the 
ashes  of  deep-sea 
weeds,  which  are  col- 
lected on  the  coasts  of  Ireland,  Scotland  and  France.  The 
percentage  of  iodin  in  seaweed  is,  however,  very  small. 

Extraction  from  Seaweed.  The  seaweed  is  first  dried  and 
then  carbonized.  The  carbonized  material  is  next  treated 
with  water,  which  dissolves  the  sodium  iodid  in  connection 


A,  Leaden  still.    B,  Earthenware  receivers. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         169 

with  some  chlorids,  bromids,  and  other  soluble  salts.  These 
are  crystallized  out  as  much  as  possible.  To  the  remaining 
liquid,  sulfuric  acid  is  added,  which  decomposes  any  carbon- 
ates or  sulfids  that  may  be  present,  and  at  the  same  time 
liberates  some  bromin  and  iodin.  The  liquid  is  then  placed 
in  a  leaden  still  (see  ill.  p.  168),  which  is  connected  with  a 
series  of  earthenware  receivers.  The  still  is  heated  gently, 
and  manganese  dioxid  is  added  from  time  to  time.  The 
vapors  of  iodin  distill  over,  are  collected,  and  are  purified 
by  redistillation. 

The  reaction  is 
2  Na  I  +  3  H2  S  04  +  Mn  02  =  2  Na  H  S  04  +  Mn  S  04  + 12  +  2  H2  0 

Industrial  Applications.  Iodin  is  of  great  use  in  medicine, 
and  in  the  manufacture  of  aniline  dyes.  It  is  also  used 
extensively  in  the  preparation  of  numerous  organic  com- 
pounds. 

Hydriodic  Add.  Hydriodic  acid  can  be  readily  prepared 
in  a  manner  similar  to  that  noted  under  hydrobromic  acid. 
The  acid  itself  is  unimportant,  but  its  salts,  the  iodids,  are 
well-known  commercial  compounds.  Potassium  iodid,  its 
principal  compound,  is  used  in  photography,  in  medicine, 
and  as  a  reagent  in  the  laboratory. 

FLUORIN 

History  and  Occurrence.  On  account  of  its  intense  affinity 
for  other  elements,  fluorin  was  not  isolated  until  1886,  when 
Moissan  succeeded  in  obtaining  it  and  in  studying  its  prop- 
erties. Fluorin  occurs  in  the  mineral  cryolite  (AlFg^NaF), 
and  as  calcium  fluorid  (CaF)  in  fluor-spar.  It  also  occurs  in 
small  quantities  in  the  bones  and  teeth  of  animals. 

Hydrofluoric  Acid.  Hydrofluoric  acid  is  made  on  a  large 
scale  in  the  same  manner  as  in  Exp.  31  £,  Part  I.,  except 


170      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


that  a  leaden  retort  is  used,  containing  the  mixed  calcium 
fluorid  and  sulfuric  acid.  The  fumes  that  are  evolved  are 
collected  in  a  tube  immersed  in  a  freezing  mixture.  This 
form  of  the  acid  can  be  stored  in  dilute  form  in  gutta-percha 
or  paraffin  bottles ;  but  if  the  acid  is  anhydrous,  it  can  be 
stored  only  in  platinum  bottles,  and  then  only  at  a  tempera- 
ture below  15°  C. 

SULFUR 

History  and  Occurrence,     From   earliest  times  sulfur  has 
been    known    to     mankind.       The 
ancient  alchemists   believed   that  it 
was  the  principle  of  combustibility. 
It  occurs  free  in  the  neighborhood 
of     volcanoes,     es- 
pecially   in    Sicily, 
where  most  of  the 
sulfur  of  commerce 
is     obtained.       In 
combination     with 
other    elements,    it 
is     found     in     the 
gases     emanating 
from    volcanoes. 
It  is  also  found  in 
the    sulf  ates  *   and 
sulfids,t    and    in 

A,  Iron  pot  containing   molten   sulfur.    B,  Pipe      organic       CO  in- 
leading  to  cylinder  C.    D,  Brick  chamber.     F,       pounds. 
Molten  sulfur.    E,  Receiver. 

*  The  most  important  are  gypsum  (calcium  sulfate,  CaSO4  2  H2O), 
and  heavy  spar  (barium  sulfate,  BaSO^. 

t  The  most  important  are  iron  pyrites  (iron  persulfid,  FeS2),  galen- 
ite  lead  sulfid  (PbS),  sphalerite  (ZnS),  and  cinnabar  (HgS). 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          I/I 

Extraction.  The  free  sulfur,  as  obtained  in  the  native 
state,  contains  earthy  impurities,  stones,  etc.  These  are 
removed  at  the  mines  by  piling  the  sulfur  in  heaps,  and  set- 
ting fire  to  it  at  the  bottom.  The  heat  of  the  burning  sulfur 
melts  the  rest  of  it,  which  runs  to  the  bottom  and  is  collected. 
This  is  shipped  away  to  other  countries,  where  it  is  further 
purified.  The  last  purification  is  done  by  melting  the  sulfur 
in  an  iron  pot  from  which  it  is  allowed  to  run  into  a  cast 
iron  retort.  The  retort  is  heated,  and  the  sulfur  vapors  'are 
led  into  a  brick  chamber.  The  first  sulfur  that  enters  the 
chamber  condenses  to  the  solid  form  as  a  fine  powder  known 
in  commerce  as  flowers  of  sulfur.  When  the  walls  of  the 
chamber  become  hot,  the  fumes  condense  to  a  liquid,  which 
is  drawn  out  into  wooden  molds,  and  when  cool,  forms  the 
roll  brimstone  of  commerce.  (See  ill.  p.  170.) 


HYDROGEN    SULFID 

History  and  Occurrence.  In  1777,  Scheele  obtained  this 
gas  by  heating  sulfur  in  hydrogen  gas.  It  is  found  in 
nature  in  volcanic  gases,  in  certain  springs,  and  wherever 
organic  matter  decomposes. 

Uses.  Hydrogen  sulfid  is  used  mainly  in  the  chemical 
laboratory  as  a  reducing  agent  and  to  precipitate  certain 
metals  as  insoluble  sulfids.*  It  is  also  extensively  used 
medicinally  wherever  springs  are  found  whose  waters  con- 
tain it  in  solution. 

The  presence  of  hydrogen  sulfid  in  eggs,  in  vulcanized 
rubber,  and  in  many  other  organic  substances  containing 

*  The  presence  of  hydrogen  sulfid  in  water  or  in  gases  can  be  de- 
termined by  means  of  the  precipitates  formed  with  solutions  of  the 
soluble  salts  of  various  metals. 


1/2      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

sulfur,  may  be  shown  by  its  blackening  a  silver  coin.  This 
is  due  to  the  breaking  up  of  the  hydrogen  sulfid,  which,  de- 
composing very  easily,  yields  nascent  sulfur.  The  black 
spot  on  the  silver  is  silver  sulfid. 


SULFUR    DIOXID 

History  and  Occurrence.  In  the  writings  of  Homer, 
mention  is  made  of  using  burning  sulfur  for  fumigation 
purposes ;  and  Pliny  tells  of  its  use  for  purifying  cloth. 
Priestley,  however,  first  prepared  pure  sulfur  dioxid,  in  1775. 
Like  hydrogen  sulfid,  sulfur  dioxid  occurs  in  nature  in  vol- 
canic gases. 

Manufacture  and  Industrial  Applications.  Sulfur  dioxid  can 
be  economically  manufactured  in  three  ways,  —  by  heating 
sulfuric  acid  and  charcoal,  by  burning  sulfur  in  a  furnace, 
and  by  roasting  iron  pyrites  (FeS2).  In  the  manufacture  of 
sulfuric  acid,  where  sulfur  dioxid  finds  its  greatest  use,  the 
last  two  methods  prevail.  Sulfur  dioxid  is  valuable  also  on 
account  of  its  bleaching  properties.  Substances  that  would 
be  injured  by  chlorin,  such  as  silk,  wool,  straw,  etc.,  are 
bleached  by  this  gas.  The  material  is  moistened  and  hung 
or  placed  in  chambers  into  which  the  fumes  of  burning  sul- 
fur are  allowed  to  enter.  Colors  acted  upon  by  chlorin  can- 
not be  restored,  but  colors  acted  upon  by  sulfur  dioxid  can 
be  restored  by  treating  them  with  dilute  sulfuric  acid  or 
with  alkalis.  Sulfur  dioxid  is  used  also  to  remove  chlorin 
from  freshly  bleached  goods.  The  sulfur  dioxid  and  the 
chlorin  here  form  sulfuric  and  hydrochloric  acids,  which  are 
afterwards  removed  by  washing.  Usually  sodium  sulfite, 
the  sodium  salt  of  sulfurous  acid,  is  used  instead  of  sulfur 
dioxid.  Sulfur  is  often  burned  in  sick  rooms,  since  the  re- 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          1/3 

suiting  sulfur  dioxid  acts  as  a  disinfectant.  Some  authori- 
ties claim,  however,  that  as  a  disinfectant  its  value  is  very 
limited. 

SULFURIC    ACID 

History  and  Occurrence.  Basil  Valentine  first  described 
sulfuric  acid  in  the  i5th  century,  though  it  was  probably 
known  before  that  time.  The  acid  was  first  made  from 
ferrous  sulfate  (FeSO4,7H2,O),  which,  when  heated,  first 
gives  off  water  and  then  sulfur  trioxid  (See  Exp.  9  e,  Part  I.). 
These  in  turn  unite,  forming  an  acid  called  fuming  sulfuric 
acid.  In  1770,  Roebuck,  of  Birmingham,  proposed  the  lead- 
chamber  process,  which  is  still  used. 

Sulfuric  acid  is  found  in  the  free  state  in  some  rivers  and 
mineral  springs. 

Manufacture.  In  Exp.  9  /,  Part  I.,  it  was  found  that  sul- 
furic acid  could  be  made  by  adding  water  to  sulfur  trioxid. 
H2  0  +  S  03  =  H2  S  04 

On  a  large  scale,  however,  this  method  is  impracticable, 
and  a  modification  of  the  method,  described  in  Exp.  39  b, 
Part  I.,  is  used.  In  Exp.  39  b,  Part  I.,  the  reaction  that 

takes  place  is 

H2  0  +  S  02  +  N  02  =  H2  S  0*  +N  0 

The  NO2  is  formed  by  the  oxidation  of  the  NO  obtained 
from  the  flask.  Thus  NO  acts  only  as  a  carrier  of  oxygen. 

Sulfuric  acid  is  manufactured  in  a  series  of  lead-lined 
chambers.  The  lead  is  used  because  it  is  the  only  inex- 
pensive metal  that  sulfuric  acid,  below  a  certain  strength, 
does  not  attack.  In  this  process  (called  the  lead-chamber 
process},  the  sulfur  dioxid  is  obtained  by  burning  sulfur  or 
by  roasting  iron  pyrites  in  a  furnace.  The  latter  method  is 

usually  used. 

2  Fe  S2  +  1 1   0  =  Fe2  03  +  4  S  02 


1/4      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

The  mixture  of  sulfur  dioxid  with  air,  both  obtained 
from  the  furnace,  passes  through  a  flue  into  a  tower  con- 
taining loosely  packed  coke.  This  is  called  a  Glover's  tower. 
From  the  Glover's  tower  the  mixed  gases  pass  into  a  lead- 
lined  chamber  into  which  enters  a  jet  of  steam ;  and  thence 
they  pass  to  a  second  similar  chamber  into  which  nitric  acid 
flows  over  a  number  of  corrugated  porcelain  cones.  These 
are  used  to  spread  the  nitric  acid  over  as  much  surface  as 
possible.  Here  a  complicated  reaction  ensues,  which,  with- 
out going  far  from  the  truth,  may  be  stated  briefly 

S  02  +  2  H  N  03  =  H2  S  04  +  2  N  02 

The  sulfuric  acid  thus  formed  contains  unused  nitric  acid 
and  oxids  of  nitrogen.  The  liquid  is  allowed  to  flow  back 
into  the  first  chamber,  where  it  meets  a  fresh  supply  of  sul- 
fur dioxid,  steam,  and  air.  The  gaseous  products  (SO2,  NO2, 
and  steam,  not  yet  combined  in  the  second  chamber)  then 
pass  into  other  lead  chambers  into  which  steam  jets  enter. 
Here  the  following  reaction  takes  place  :  — 

S  02  +  N  02  -f-  H2  0  =  H2  S  04  +  N  0 
The  NO  is  also  oxidized  to  NO2  and  N2O3, 

8  N  0  +  3  02  -  2  N3  03  +  4  N  02 

The  last  of  these  chambers  is  connected  with  a  tower, 
called  a  Gay  Lussac  tower,  filled  with  loosely  packed  coke 
through  which  slowly  trickles  a  stream  of  strong  sulfuric 
acid.  This  acid  dissolves  the  unused  oxids  of  nitrogen  that 
would  otherwise  escape  with  the  waste  nitrogen  left  from  the 
air,  and  is  then  collected  at  the  bottom  and  forced  to  the 
top  of  the  Glover's  tower  by  means  of  compressed  air.  It 
there  trickles  down  through  the  coke,  and  meets  a  fresh 
supply  of  su|fur  dioxid  and  air,  thus  making  more  sulfuric 
acid. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS 


AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

The  sulfuric  acid  thus  obtained  is  about  65  %  pure,  and 
has  a  specific  gravity  of  about  1.5.  It  is  removed  to  leaden 
pans,  and  evaporated  until  it  reaches  a  specific  gravity  of 
1.75.  After  it  reaches  this  strength,  lead  cannot  be  used, 
because  it  is  attacked  by  the  acid.  Further  concentration 
is  reached  by  evaporating  it  in  platinum  dishes  until  the  acid 
reaches  a  specific  gravity  of  1.84.  Instead  of  platinum, 
iron  vessels  have  recently  been  used,  since  iron  is  not  at- 
tacked by  the  concentrated  acid. 

Industrial  Applications.  Sulfuric  acid  is  the  most  widely 
used  of  all  chemicals.  Thousands  of  tons  are  produced 
annually.  It  is  a  necessity  in  the  manufacture  of  most  acids 
and  salts.  The  manufacturers  of  soda  ash,  glass,  soap,  and 
fertilizers,  are  great  consumers  of  sulfuric  acid,  in  fact,  there 
is  scarcely  a  manufacturing  plant  of  any  kind  that  does  not 
use  the  acid  for  some  purpose  or  other. 

SELENIUM    AND    TP:LLURIUM 

History  and  Occurrence.  There  are  two  elements,  selenium 
and  tellurium,  that  are  very  similar  to  sulfur  both  in  their 
chemical  properties  and  in  their  compounds.  Both  are  rare 
elements.  Selenium  generally  occurs  combined  with  copper, 
lead,  or  silver.  It  is  a  dark,  reddish-brown,  crystalline  solid 
in  one  of  its  allotropic  forms,  while  in  the  other  it  is  a  dark, 
gray  solid.  Tellurium  usually  occurs  as  a  tellurid  of  silver, 
gold,  or  bismuth,  and  is  a  white,  brittle  solid  of  metallic 
luster.  The  former  element  was  discovered  by  Berzelius  in 
1817,  the  latter  by  Klaproth  in  1798. 

NITROGEN 

History.  In  1772,  Rutherford  showed  that  if  animals 
breathe  in  a  confined  volume  of  air,  and  the  carbon  dioxid 


HISTORY,     OCCURRENCE    AND    APPLICATIONS 

thus  formed  is  removed  by  means  of  lime  water,  a  gas  is  left 
that  does  not  support  combustion.  In  the  same  year, 
Priestley  showed  that  the  same  thing  is  true  of  confined  air 
in  which  carbon  has  been  burnt.  Lavoisier  gave  it  the 
name  of  azote,  but  the  name  nitrogen,  suggested  by  Chaptal, 
was  generally  accepted. 

Occurrence.  About  four-fifths  of  the  air  is  nitrogen  (see 
Exp.  36,  Part  I.).  This  gas  is  also  a  constituent  of  a  large 
number  of  organic  compounds,  and  occurs  in  combination 
with  hydrogen  as  ammonia,  and  with  metals  and  oxygen  as 
nitrates. 

AMMONIA 

History.  The  early  alchemists  were  familiar  with  the  odor 
of  the  salt  now  called  ammonium  carbonate.  Basil  Valen- 
tine showed  that,  when  an  alkali  acts  upon  sal-ammoniac, 
a  strong  smelling  solution  is  obtained.  The  name  sal- 
ammoniac  was  given  to  a  salt  that,  in  the  seventh  century, 
was  brought  to  Europe  from  Asia.  It  was  later  obtained  by 
the  dry  distillation  of  animal  matter,  such  as  hoof,  hair, 
horn.  etc.  The  carbonate  of  ammonium  obtained  in  this 
way  was  neutralized  with  hydrochloric  acid,  and  ammonia 
was  obtained  from  the  resulting  ammonium  chlorid.  In 
1774,  Priestley  heated  sal-ammoniac  with  lime,  and  collected 
the  resulting  gas  over  mercury.  By  passing  electric  sparks 
through  the  gas,  he  found  it  to  be  a  compound  of  hydrogen 
and  nitrogen. 

Occurrence.  Ammonia  occurs  in  small  quantities  in  the 
air,  where  it  is  believed  to  be  formed  by  electric  discharges. 
It  is  formed  also  by  the  decay  of  animal  matter.  Ammonium 
salts  occur  in  the  soil  and  in  certain  mineral  waters. 

Manufacture  and  Applications.  Practically  all  the  ammonia 
of  commerce  is  obtained  as  a  by-product  in  the  manufacture 


AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

of  coal  gas.  When  coal  is  heated  in  the  absence  of  air,  the 
nitrogen  passes  off  in  combination  with  hydrogen  as  ammonia. 
This  dissolves  in  the  condensed  moisture  which  is  formed 
at  the  same  time.  From  this  liquid,  ammonia  is  obtained. 

Ammonia  has  come  to  be  one  of  the  most  important  of 
commercial  chemicals.  This  has  been  due  in  a  great  meas- 
ure to  its  use  in  making  artificial  ice.  Ammonia  gas  con- 
denses to  a  liquid  under  a  pressure  of  seven  atmospheres  at 
the  ordinary  temperature.  When  it  is  allowed  to  evaporate, 
a  great  amount  of  heat  is  rendered  latent,  hence  its  value  in 
refrigeration.  It  is  also  used  in  the  preparation  of  a  great 
number  of  compounds,  some  of  which  will  be  mentioned 
later.  The  value  of  compost  as  a  fertilizer  is  due  to  the 
ammonia  evolved  by  the  decomposition  of  the  animal  matter 
in  it. 

OXIDS    OF   NITROGEN 

Nitrogen  forms  five  oxids  as  follows  :  — 

Nitrous  oxid,  N2  0     (See  Exp.  46,  Part  I.) 
Nitric  oxid,  N  0     (See  Exp.  39,  Part  I.) 
Nitrogen  trioxid,  N2  Os 
Nitrogen  peroxid,  N  02 
Nitrogen  pentoxid,  N2  Os 

Only  two  are  important,  the  nitrous  oxid  and  the  nitric 
oxid.  The  former  is  the  well-known  laughing  gas,  used  by 
the  dentist ;  while  the  latter  finds  its  chief  use  in  the  manu- 
facture of  sulfuric  acid. 

ACIDS    OF   NITROGEN- 

Of  the  three  acids  of  nitrogen,  viz., 

Hyponitrous  acid,  H  N  0 
Nitrous  acid,  H  N  02 
Nitric  acid,  H  N  03 


HISTORY,    OCCURRENCE   AND  APPLICATIONS 


only  nitric  acid  is  important.     Some  salts  of  the  other  two, 
however,  are  extensively  used. 


NITRIC  ACID 

History.  Geber,  the  Arabian  alchemist,  seems  to  have 
been  the  first  person  to  make  nitric  acid.  It  was  called  aqua 
fortis  by  the  alchemists.  The  present  method  of  making 
nitric  acid  from  a  nitrate  and  sulfuric  acid  was  first  spoken 
of  in  the  writings  of  Glauber  in  the  iyth  century.  Caven- 
dish, in  1785,  determined  definitely  that  it  was  a  compound 
of  hydrogen,  nitrogen,  and  oxygen. 


A,  Retort.    B,  Earthenware  receivers. 

Occurrence.  Nitric  acid  is  probably  found  in  the  atmos- 
phere, since  we  might  expect  it  to  be  formed  there  by  electric 
discharges.  It  occurs  plentifully  on  the  earth  in  the  form 
of  nitrates,  the  chief  of  which  are  potassium  nitrate  (saltpeter) 
and  sodium  nitrate  (Chili  saltpeter).  Whenever  nitrogenous 
organic  matter  decomposes  in  the  air,  the  nitrogen  takes  the 
form  of  ammonia.  If  an  alkali  is  present,  a  slow  oxidation 
takes  place,  and  forms  a  nitrate  of  the  alkali  present.  In 
this  manner,  the  nitrate  beds  of  India,  Chili,  and  Peru  have 
been  formed. 


ISO      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Manufacture.  Nitric  acid  is  made  for  commercial  use  in 
the  same  way  as  in  the  laboratory  (see  Exp.  37,  Part  I.). 
Large  cast-iron  retorts,  protected  inside  by  a  lining  of  clay, 
are  used  in  which  sodium  nitrate  and  sulfuric  acids  are 
placed.  The  nitric  acid  formed  distills  over  into  earthen- 
ware bottles.  (See  ill.  p.  179.)  It  is  then  purified  by  redis- 
tillation with  an  equal  volume  of  sulfuric  acid.  It  is  then 
rendered  colorless  by  means  of  a  current  of  dry  air  made  to 
pass  over  it  while  it  is  gently  warming.  In  some  establish- 
ments, instead  of  the  earthenware  bottles,  a  system  of  tubes 
called  "  Hart's  condensation  tubes  "  are  used,  in  which  the 
acid  is  condensed  and  purified. 

Industrial  Applications.  As  has  been  stated  before,  nitric 
acid  is  extensively  used  in  the  manufacture  of  sulfuric  acid, 
and,  in  that  process,  is  the  source  of  the  necessary  nitric  oxid. 
Another  application  of  nitric  acid  is  in  the  manufacture  of 
high  explosives  such  as  nitroglycerin,*  gun  cotton, f  etc. 
It  is  used  also  in  the  preparation  of  metallic  nitrates,  and  in 
the  manufacture  of  many  dyes.  Its  salt,  potassium  nitrate, 
is  used  in  the  manufacture  of  gunpowder. 

*  When  glycerin,  C3H,(OH)3,  is  treated  with  a  mixture  of  concen- 
trated sulfuric  and  nitric  acids  at  the  proper  temperature,  a  yellow  oil  is 
obtained.  This  oil  is  called  nitroglycerin.  It  is  very  unstable,  and  is 
especially  sensitive  to  concussion.  It  may  be  burned  in  the  open  air 
but  explodes  if  heated  too  highly.  Dynamite  is  a  certain  peculiar  kind 
of  earth  saturated  with  nitroglycerin,  and  can  be  handled  with  compara- 
tive safety  by  persons  understanding  its  use. 

The  action  of  the  acid  on  the  glycerin  may  be  represented  by 

C3Hs(OH)3  +  3  HN03  =  C3H5(N03)3  +  3  H2  0 
The  sulfu-ic  acid  combines  with  the  water  formed  in  the  operation. 

t  Gun  cotton  is  made  by  treating  cotton  fiber  with  a  mixture  of  con- 
centrated sulfuric  and  nitric  acids.  The  reaction  is  similar  to  that  of 
nitroglycerin. 


HISTORY,    OCCURRENCE   AND    APPLICATIONS 


ARGON 

History.  Argon  was  discovered  in  the  air  in  1894  by 
Rayleigh  and  Ramsey.  It  was  found  that  nitrogen  obtained 
from  the  air,  and  nitrogen  obtained  from  chemicals  did  not 
weigh  the  same.  On  account  of  this  fact,  the  suspicion  arose 
that  there  might  be  some  hitherto  unknown  element  or 
elements  present  in  the  air.  After  long  experimentation, 
the  water,  carbon  dioxid,  oxygen,  and  nitrogen  were  removed 
from  air,  and  a  new,  inert,  colorless  gas,  having  a  density  of 
20,  was  obtained.  To  obtain  the  new  gas,  the  moisture  and 
carbon  dioxid  were  removed  from  the  air,  after  which  the 
oxygen  was  removed  by  passing  the  dry  air  over  red-hot 
copper.  The  nitrogen  was  then  removed  by  passing  the 
remainder  through  a  tube  containing  magnesium  turnings 
heated  white-hot. 

It  has  since  been  found  by  the  same  experimenters  that 
the  gas  obtained  by  the  above  process  is  not  a  single  gas, 
but  that  there  is  present  probably  a  number  of  different 
gases,  very  similar  in  their  properties.  Of  these%  neon, 
metargon,  and  xenon  have  been  separated. 

HELIUM 

While  investigating  argon,  Rayleigh  discovered  in  certain 
rare  minerals  an  element  that  had  hitherto  been  known  only 
to  exist  in  the  sun.  It  had  already  been  called  helium.  Its 
atomic  weight  was  found  to  be  about  4. 

PHOSPHORUS 

History  and  Occurrence.  Brand  claimed  to  have  discovered 
phosphorus  in  1669.  From  its  property  of  becoming  lumi- 
nous in  the  dark,  its  name  is  taken  from  the  words  <£u>s,  light, 


1 82      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

and  (j>epa>,  I  bear.  The  element  never  occurs  free  in  nature. 
It  is  widely  distributed,  however,  in  combination.  Many 
rocks  contain  phosphorus  compounds.  When  these  rocks 
disintegrate,  the  soils  that  are  formed  become  very  produc- 
tive on  account  of  the  presence  of  phosphorus.  Vegetation, 
in  turn,  is  transformed  into  animal  matter.  Here  the  phos- 
phorus appears  mostly  in  the  bones  as  calcium  phosphate. 
A  fossil  substance  called  "  caprolites  "  also  contains  calcium 
phosphate,  and  it  is  from  this  and  from  bones  that  the 
phosphorus  of  commerce  is  derived. 

Manufacture.  In  the  manufacture  of  phosphorus  from 
bones,  the  first  step  is  to  remove  the  non-phosphorus-bear- 
ing organic  matter.  This  is  done  by  burning.  The  ash  is 
then  treated  with  sulfuric  acid,  which  changes  the  insoluble 
calcium  phosphate  to  soluble  acid  calcium  phosphate,  accord- 
ing to  the  following  reaction  : 

Ca3  (P  04)2  +  2  H2  S  04  =  Ca  H4  (P  04)2  +  2  Ca  S  04 

After  the  calcium  sulfate  has  been  removed,  the  acid  calcium 
phosphate  solution  is  evaporated  to  a  syrup  and  mixed  into 
a  paste  with  charcoal  powder.  This  is  heated  to  redness 
in  an  earthenware  retort,  whose  mouth  dips  under  water. 
Water  is  at  first  driven  off  thus :  — 

Ca  H4  (P  04)2  =  Ca  (P  03)2  +  2  H2  0 

The  metaphosphate  thus  formed  then  reacts  with  the  char- 
coal as  follows :  — 

3  Ca  (P  03)2  +  10  C  =  P4  +  Ca3  (P  04)2  +  10  C  0 

The  phosphorus  distills  over  into  the  water,  is  collected,  and 
is  purified  by  redistillation.  Sometimes  sand  is  added  to- 
gether with  charcoal,  in  which  case  all  the  phosphorus  is 
obtained. 

The  phosphorus  obtained  by  the  above  method  is  the  yel- 
low variety.  The  red  amorphous  phosphorus  is  made  by 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          183 

heating,  for  a  number  of  days,  the  yellow  variety  at  240°  C. 
in  an  iron  vessel  having  only  a  small  opening.  Not  all  the 
yellow  phosphorus  changes  over.  The  unchanged  phos- 
phorus is  dissolved  out  of  the  cooled  mixture  by  carbon 
bisulfid,  or  is  removed  by  means  of  sodium  hydroxid  which 
acts  upon  the  yellow  phosphorus  and  forms  phosphoretted 
hydrogen. 

Industrial  Applications.  Phosphorus  is  used  in  the  arts 
principally  in  the  manufacture  of  matches.  The  old  friction 
matches  were  made  by  first  dipping  the  wood  into  melted 
sulfur,  cooling  it,  and  then  tipping  it  with  a  mixture  of 
phosphorus,  glue,  and  some  oxidizing  agent.  The  so-called 
safety  matches  do  not  contain  phosphorus.  Instead  of  this 
they  are  tipped  with  antimony  sulfid.  The  box  is  painted 
with  a  preparation  of  red  amorphous  phosphorus,  antimony 
sulfid,  and  glue.  Sometimes  manganese  dioxid  or  some 
other  oxidizing  agent  is  used  in  this  mixture. 

Phosphin.  There  are  three  compounds  of  phosphorus  and 
hydrogen,  only  one  of  which,  phosphin  (PH3),  will  be  noted 
here.  If  phosphorus  is  heated  in  a  solution  of  sodium  or 
potassium  hydroxid,  a  gas  is  liberated  that  ignites  spon- 
taneously. If  the  gas  is  led  under  water,  the  bubbles  on 
reaching  the  surface  burst  into  flame  and  form  white  rings 
of  phosphorus  pentoxid.  The  gas  is  colorless,  has  a  very 
disagreeable  odor,  and  is  poisonous.  It  forms  a  class  of 
substances,  called  phosphonium  compounds,  analogous  to 
those  formed  by  ammonium. 

Phosphoric  Acid.  The  phosphoric  acid  of  commerce  is 
made  from  bones.  The  main  value  of  phosphoric  acid  lies 
in  its  salts,  which  are  the  principal  constituents  of  fertilizers. 
For  the  purpose  of  making  fertilizer,  bone  ash  is  treated 
with  sulfuric  acid  to  obtain  the  soluble  acid  calcium  phos- 
phate, CaH4(P04)2.  . 


1 84      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


ARSENIC 

History  and  Occurrence.  Arsenic  compounds  were  known 
to  the  alchemists,  who  thought  they  could  use  them  in  trans- 
muting the  metals.  Brandt,  in  1773,  first  showed  that  white 
arsenic  was  the  calx  of  a  metal.  Later,  its  position  as  an 
oxid  was  established. 

Arsenic  occurs  in  the  free  state  in  nature,  but  usually  it  is 
found  in  combination  with  iron,  cobalt,  nickel,  or  sulfur. 

Applications.  Arsenic  is  used  in  the  manufacture  of  cer- 
tain pigments.  In  the  manufacture  of  glass,  arsenic  oxid  is 
used  to  remove  the  green  tint  given  by  ferrous  hydroxid, 
which  by  its  action  becomes  ferric  oxid.  In  the  manufacture 
of  shot,  arsenic  has  the  effect  of  hardening  the  lead.  In 
medicine,  it  is  sometimes,  used  as  a  tonic.  For  preserving 
skins,  it  is  valuable  to  the  taxidermist.  In  the  form  of  Paris 
green  (CuHAsO3),  it  is  valuable  as  a  poison  for  the  de- 
struction of  insects. 

ANTIMONY 

History.  Mention  is  made  in  the  Scriptures  of  the  metal 
we  call  antimony.  In  early  times  the  sulfid  was  used  by 
the  women  of  the  East  to  paint  their  eyebrows.  Pliny  called 
it  stibium,  although  in  Latin  it  was  also  known  as  antimonium. 

Occurrence  and  Applications.  While  antimony  occurs  some- 
times free  in  nature,  it  is  usually  found  as  the  sulfid  (Stib- 
nite,  Sb2S3).  When  the  ore  is  roasted  in  the  presence  of  air, 
the  oxid  (Sb2O3)  is  formed.  This,  mixed  with  carbon  and 
ignited  strongly,  gives  the  metal. 

Antimony  is  used  chiefly  in  making  alloys.  It  hardens 
the  alloy,  and  also,  by  its  property  of  expanding  when  cooled, 
makes  it  invaluable  in  the  manufacture  of  type  metal.  Type 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          185 

metal  is  composed  of  about  two  parts  lead,  and  one  part 
each  of  antimony  and  tin.  Antimony  is  used  also  in  making 
white  metal,  pewter,  and  Britannia  metal.  Finely  divided  anti- 
mony, prepared  by  zinc  from  the  chlorid,  is  called  antimony 
black,  and  is  used  for  giving  a  metallic  appearance  to  plaster 
casts,  statues,  etc.  In  medicine  its  compound,  potassium 
antimony  tartarate  (tartar  emetic),  is  used  as  an  emetic. 

BORON 

History  and  Occurrence.  Gay  Lussac,  Thenard,  and  Sir 
Humphry  Davy  obtained  boron  in  the  elementary  state  in 
1808.  There  are  two  varieties,  the  amorphous  and  the  crys- 
talline boron. 

It  never  occurs  free,  but  is  found  as  boric  acid  (H3BO3) 
in  Tuscany,  and  as  sodium  salts  in  California  and  Thibet. 


Boric  Acid,  ffzBOz.  In  Tuscany,  jets  of  steam  contain- 
ing boric  acid  escape  to  the  surface  of  the  earth  from  sub- 
terranean sources.  Brick  basins  are  built  around  these 
steam  jets.  The  heat  causes  the  solution  obtained  to 
evaporate.  When  sufficiently  strong,  the  solution  is  removed 
and  allowed  to  crystallize. 

Uses.     Boric  acid  is  a  valuable  antiseptic,   and  for  this 


1 86       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

reason  is  used  extensively  in  surgery.     It  is  also  sometimes 
used  in  preserving  perishable  foods. 

Borax,  Na^B^O1  is  a  valuable  compound,  and  is  much 
used  as  a  flux  by  tin-smiths  and  copper-workers.  It  gives  a 
clean,  bright,  metallic  surface  on  account  of  its  property  of 
dissolved  metallic  oxids. 


CARBON 

History.  Carbon  has  of  course  been  known  from  ancient 
times,  but  its  allotropic  forms  were  not  understood  until  the 
end  of  the  i8th  century,  and  its  relation  to  organic  chemistry 
not  until  somewhat  later.  In  the  form  of  graphite,  it  was 
known  to  the  alchemists.  Graphite  pencils  were  first  made 
in  1565.  In  1772,  Lavoisier  showed  that  diamond  and 
charcoal  were  chemically  identical.  He  burned  a  diamond 
and  obtained  carbon  dioxid.  Tennant,  in  1776,  proved 
that  like  weights  of  charcoal,  graphite,  and  diamond  give 
like  weights  of  carbon  dioxid. 

Occurrence.  Carbon  occurs  as  diamond  principally  in 
India,  South  Africa,  and  Brazil.  It  is  usually  transparent 
or  slightly  tinged  with  yellow,  although  it  is  sometimes  found 
red,  green,  blue,  or  even  black.  As  graphite,  carbon  occurs 
widely  distributed,  being  found  mostly  in  England,  Siberia, 
Ceylon,  Canada,  New  York,  and  California. 

Graphite  is  used  in  the  manufacture  of  lead  pencils,  cruci- 
bles, and  various  lubricants.  It  is  also  used  in  foundries 
for  facings,  and  in  electrotyping. 

Amorphous  Carbon.  Coal.  As  mineral  coal,  carbon  is 
found  in  almost  all  countries  of  the  world.  The  most,  how- 
ever, is  found  in  England  and  the  United  States.  Coal  is 
all  that  is  left  of  the  great  primeval  forest  that  covered  the. 
earth  long  before  the  advent  of  man.  When  vegetable 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         l8/ 

matter  decays  in  the  absence  of  air,  in  the  earth,  or  under 
water,  it  gives  off  gases,  and  a  substance  similar  to  coal 
remains.  In  some  parts  of  the  earth,  the  trees  of  the  ancient 
forests  were  buried  by  earthy  material,  and  underwent  the 
process  of  decay  for  millions  of  years.  Upon  the  complete- 
ness of  this  process  depend  the  different  kinds  of  coal. 
Where  the  decomposition  has  been  most  thorough,  anthra- 
cite coal  results.  Bituminous  coal  is  rich  in  hydrocarbons, 
which  it  gives  off  when  heated.  Cannel  coal  is  a  variety  of 
bituminous  coal,  and  is  especially  rich  in  hydrocarbons. 
Brown  coal  and  peat  belong  to  a  later  geological  period  than 
the  others  mentioned. 

Charcoal.  Charcoal  is  obtained  by  the  partial  combus- 
tion of  wood.  The  wood  is  arranged  in  a  pile,  and  covered 
with  earth.  It  is  then  ignited,  and  allowed  to  burn.  On 
account  of  being  protected  from  the  air,  the  combustion  is 
imperfect,  and  only  the  volatile  part  of  the  wood  is  driven 
off. 

Animal  Charcoal.  By  heating  bones  and  other  kinds  of 
animal  refuse  in  iron  retorts,  a  variety  of  charcoal  tha.t  con- 
tains calcium  phosphate  is  obtained.  It  is  largely  used  as  a 
filter  for  removing  vegetable  coloring  matters  from  liquids. 

Coke  and  Gas  Carbon.  In  the  retorts  of  the  gas  works, 
after  the  volatile  products  have  been  driven  off  from  the 
coal,  there  remains  behind  a  gray,  porous  solid  that  is  called 
coke.  It  is  also  made  by  burning  coal  in  ovens  so  as  to 
burn  out  only  the  volatile  products.  In  the  upper  parts  of 
the  gas  retorts,  there  remains  a  dull  black  mass  known  as 
gas  retort  carbon.  It  is  used  in  making  carbon  plates  for 
electric  batteries,  and  formerly  for  making  pencils  for  arc- 
light  lamps.  Gas  retort  carbon  is  the  purest  form  of  amor- 
phous carbon. 

Lamp  black  is  made  by  burning  turpentine  and  other  oils 


1 88      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

rich  in  hydrocarbons,  and  collecting  the  smoke.  It  is  used 
chiefly  in  printer's  ink  and  black  paint. 

Hydrocarbons.  In  Pennsylvania,  in  Ohio,  and  in  the  Cau- 
casus, an  oily  liquid  called  petroleum  is  found  in  the  earth. 
This  liquid  is  composed  mostly  of  a  mixture  of  hydrocar- 
bons. These  are  partly  gaseous,  partly  liquid,  and  partly 
solid.  The  petroleum  is  distilled,  and  the  resulting  mixtures 
are  afterwards  washed  with  water  and  alkali.  Gasolene, 
naphtha,  benzine,  kerosene,  and  paraffin,  are  among  the 
principal  substances  obtained.  These  are  all  mixtures  of 
various  hydrocarbons. 

Natural  gas  occurs  in  the  earth  in  large  quantities,  usually 
near  coal  beds.  Its  formation  is  supposed  to  be  due  to  the 
dry  distillation  of  coal  in  the  interior  of  the  earth. 

The  hydrocarbons  form  a  definite  series  of  compounds, 
each  hydrocarbon  of  the  series  differing  from  the  next  by  an 
atom  of  carbon  and  two  atoms  of  hydrogen,  i.e.,  by  CH2. 
The  first  hydrocarbons  of  the  simplest  groups  are 

Methane,  C  H4  (See  Exp.  33  e.  Part  I.) 
Ethylene,  C2  H4  (See  Exp.  33 /,  Part  I.) 
Actylene,  C2  H2 

METHANE    OR    MARSH    GAS 

History  and  Occurrence.  This  gas  is  mentioned  by  Pliny, 
and  Basil  Valentine  speaks  of  its  presence  in  mines.  In 
1785,  Berthollet  proved  that  methane  contained  both  hydro- 
gen and  carbon,  and  in  1805  Henry  showed  the  difference 
between  methane  and  ethylene. 

Methane  occurs  in  nature  in  mines,  and  wherever  vege- 
table matter  is  decaying  under  water ;  hence  its  common 
name,  marsh  gas.  Miners  call  it  fire  damp.  Natural  gas  is 
rich  in  methane. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          189 


ETHYLENE 

History  and  Occurrence.  In  the  lyth  century,  Becher  dis- 
covered ethylene  by  heating  alcohol  with  sulfuric  acid,  but  it 
was  not  until  Henry  took  up  the  study  of  methane  and 
ethylene  that  its  composition  was  understood.  It  occurs  in 
the  gases  that  emanate  from  oil  wells,  and  is  the  principal 
constituent  of  coal  gas. 

ACETYLENE 

History  and  Application.  Acetylene  was  discovered  by 
Edmund  Davy  in  1836.  It  is  formed  by  the  incomplete 
combustion  of  other  hydrocarbons,  such  as  ethylene,  coal 
gas,  etc.  For  instance,  when  the  flame  "  backs  down  "  in  a 
Bunsen  burner,  acetylene  is  formed. 

Acetylene  is  rich  in  carbon,  burns  with  a  bright,  white 
flame,  and,  if  it  can  be  produced  cheaply  enough  and  regu- 
lated with  perfect  safety,  is  likely  to  become  one  of  the 
chief  sources  of  illumination. 

It  is  now  made  by  the  action  of  calcium  carbid  \^CaC2) 
upon  water. 

Ca  C2  +  2  H2  0  =  C2  H2  +  Ca  (0  H)2 

ILLUMINATING  GAS 

William  Murdock,  a  Scotchman,  first  saw  the  practicabil- 
ity of  making  illuminating  gas  from  coal.  This  was  in  1792, 
but  London  was  not  lighted  by  gas  until  1812,  nor  Paris 
until  1815. 

When  coal  is  heated  in  the  absence  of  air,  three  classes 
of  products  are  formed, 

1.  Illuminating  gas. 

2.  Coal  tar,  a  thick,  oily,  strong-smelling  liquid. 


AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


3.  Ammoniacal 
liquor,  containing 
ammonia  and  various 
other  compounds  in 
solution. 

When  the  coal  is 
heated  in  the  gas 
retorts,  the  three 
classes  of  products 
above  mentioned 
pass  through  a  pipe 
that  dips  below  water 
in  another  large  pipe 
called  the  hydraulic 
main.  The  water 
absorbs  some  of  the 
soluble  gases,  and 
takes  up  some  of  the 
ammoniacal  liquor 
and  tarry  products. 
From  the  hydraulic 
main,  the  gas  passes 
through  a  series  of 
upright  iron  pipes 
beneath  which  is 
water.  As  the  gas 
passes  through  these 
pipes,  it  cools,  and 
another  portion  of 
the  tarry  products 
and  ammoniacal  liq- 
uor is  deposited, 
which  runs  down  into 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         IQI 

the  water.  From  this  pipe,  the  gas  passes  to  a  series  of  towers, 
called  scrubbers,  filled  with  coke.  It  passes  into  the  bottom  of 
one  tower,  through  the  coke,  and  then  at  the  top  meets  a 
spray  of  ammonia  water.  In  like  manner,  it  passes  through 
the  next  tower,  and  so  on.  The  ammonia  water  removes 
most  of  the  hydrogen  sulfid.  There  is  still  left  in  the  gas 
some  hydrogen  sulfid,  carbon  bisulfid,  and  carbon  dioxid.  To 
remove  these,  the  gas  passes  through  a  series  of  chambers 
containing  lime  or  hydrated  ferric  oxid.  The  gas  is  then 
ready  for  use,  and  passes  into  the  gas  holder. 


CARBON    MONOXID 

History  and  Occurrence.  This  compound  (see  Exp.  33, 
Part  I.)  was  studied  by  various  investigators  during  the  last 
part  of  the  i8th  century,  but  it  was  some  time  before  its 
true  character  was  explained.  It  does  not  occur  free  in 
nature,  but  it  is"  formed  wherever  carbon  burns  in  an  in- 
sufficient supply  of  air.  \ 

Applications.  The  principal  applications  of  carbon  mon- 
oxid  are  in  the  manufacture  of  water  gas  and  in  the  reduc- 
tion of  metals  from  their  ores. 


WATER   GAS 

In  the  manufacture  of  water  gas  superheated  steam  is 
passed  over  red  hot  anthracite  coal,  giving  the  gases  carbon 
monoxid  and  hydrogen. 


Both  of  these  gases  burn  with  a  blue  flame,  hence  they  must 
be  enriched  with  some  gas  that  burns  brightly. 


192      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


CARBON    DIOXID 

History  and  Occurrence.  Van  Helmont,  one  of  the  early 
investigators,  held,  as  early  as  the  beginning  of  the  iyth 
century,  that  this  gas  was  different  from  air.  He  called  it 
"gas  sylvestre."  Its  chemical  nature  was  first  determined 
by  Lavoisier. 

Carbon  dioxid  occurs  free  in  the  air  and  in  combination 
with  metals  as  carbonates.  It  is  also  found  in  some  mineral 
waters. 

Manufacture  and  Applications.  For  commercial  purposes, 
carbon  dioxid  is  made  by  treating  sodium  carbonate  or  cal- 
cium carbonate  with  acids.  Under  38.5  atmospheres  pres- 
sure and  at  o°  C.,  this  gas  condenses  to  a  liquid.  It  is 
then  stored  in  steel  cylinders,  from  which  it  can  be  taken  at 
will.  Its  principal  use  is  in  the  manufacture  of  aerated 
beverages. 

CARBON    BISULFII) 

This  compound  (CS2)  does  not  occur  in  nature,  but  is 
made  by  leading  vapors  of  sulfur  over  red-hot  charcoal.  It 
is  extensively  used  in  the  arts  as  a  solvent  for  rubber,  phos- 
phorus, sulfur,  iodin,  and  many  oils  and  gums. 

CYANOGEN 

By  heating  mercuric  cyan  id,  a  colorless  gas  is  obtained 
that  burns  with  a  beautiful  purple  flame,  and  has  the  odor  of 
peach  kernels. 

Hg  (C  N),  =  (C  N)2  +  Hg 

It  combines  with  hydrogen,  forming  hydrocyanic  acid  (HCN), 
commonly  called  prussic  acid.  This  is  such  a  deadly  poison 
that  it  is  used  only  in  dilute  solutions. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS          193 

SILICON 

History  and  Preparation.  Silica,  or  sand,  was  believed  to 
be  a  compound  body  long  before  Berzelius  first  obtained 
impure  silicon  by  fusing  together  iron,  carbon,  and  silica. 
It  is  best  prepared  by  heating  potassium  silico-fluorid  and 
metallic  potassium  in  an  iron  tube. 

K2  Si  F6  +  K4  =  6  K  F  +  Si 

Silicon  is  a  brown  powder  that,  when  heated  in  the  air, 
burns  to  silicon  dioxid.  If  this  amorphous  silicon  is  fused 
with  zinc,  it  forms  dark  glittering  crystals  of  silicon  that  may- 
be obtained  by  dissolving  away  the  zinc  with  an  acid. 

Occurrence.  Next  to  oxygen,  silicon  is  the  most  abundant 
element  in  nature.  It  occurs  in  combination  with  oxygen  in 
quartz,  and  in  the  form  of  silicates  of  the  metals. 

SILICON    DIOXID 

Quartz  is  the  purest  form  of  silicon  dioxid.  Sand  and 
sandstone  are  other  forms.  Silicon  dioxid  is  used  in  the 
arts  in  the  manufacture  of  glass,  porcelain,  and  in  potter)-. 


GLASS 

The  Egyptians  are  believed  to  be  the  first  people  that 
manufactured  glass.  On  Egyptian  tombs  are  found  pictures 
of  glass  blowers  carrying  on  their  vocation.  During  the 
Middle  Ages,  Venice  was  famous  for  its  glass  manufactures ; 
but  after  their  decay  the  art  passed  to  the  workmen  of 
Bohemia.  The  art  of  glass  making  has  steadily  improved, 
until  at  present  some  of  our  most  beautiful  and  marvelous 
works  of  art  are  due  to  the  glass-worker's  skill.  Glass  is  a 


IQ4      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

mixture   of  various   silicates,   especially  of  the   alkalis   and 
alkaline  earths. 

The  various  kinds  of  glass  are  divided  into  classes  accord- 
ing to  their  ingredients. 

1.  Bohemian  glass. 

2.  Crown  or  window  glass.  , 

3.  Common  green  or  bottle  glass. 

4.  Flint  glass  or  crystal. 

Bohe?nian  glass  is  a  silicate  of  potash  and  lime,  is  fusible 
with  difficulty,  and  withstands  chemical  reagents  better  than 
any  other  kind. 

Crown  glass  is  a  silicate  of  soda  and  lime.  It  is  more 
readily  fusible,  and  is  more  easily  acted  upon  by  chemicals 
than  is  Bohemian  glass. 

Bottle  glass  is  a  silicate  of  soda  and  lime,  mixed  with  the 
oxids  of  aluminum  and  iron.  The  green  color  is  due  to  the 
iron.  This  color  varies  from  green  to  brown,  depending 
upon  impurities. 

Fhnt  glass  is  potash-lead  silicate.  It  is  the  softest  kind  of 
glass,  and  has  a  bright  luster  and  high  refractive  power. 
On  this  account,  it  is  used  in  making  lenses  for  optical  in- 
struments. 

Glass  can  be  colored  by  means  of  various  metallic  oxids. 
Gold  compounds  give  it  a  beautiful  ruby  tint.  Cuprous  oxid 
colors  glass  an  intense  red,  while  cupric  oxid  colors  it  green. 
Cobalt  gives  blue,  manganese  gives  violet.  Black  is  obtained 
by  the  addition  of  sesqui-oxid  of  iridium. 

PORCELAIN 

Porcelain  differs  from  glass  in  that  it  is  made  of  kaolin, 
a  silicate  of  aluminum,  Al2(SiO3)3.  Porcelain  is  glazed 
after  it  has  been  "  fired  "  for  the  first  time.  By  one  method, 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         1 95 

the  ware  is  dipped  into  a  glazing  material,  usually  finely  pow- 
dered quartz  and  feldspar,  is  then  dried,  placed  in  earthen- 
ware pots,  and  heated  in  a  furnace.  The  glaze  fuses  and 
spreads  over  the  surface  of  the  ware.  The  ware  is  then 
allowed  to  cool  in  the  furnace  to  anneal  it;  otherwise  it 
would  be  brittle  on  account  of  the  unequal  tension  of  its  dif- 
ferent parts.  Ordinary  pottery  is  simply  baked,  and  not 
glazed.  Bricks  are  simply  baked  clay.  The  red  color  is 
due  to  the  presence  of  iron  silicate. 


POTASSIUM 

History  and  Occurrence.  Until  1807,  the  alkalis  were 
believed  to  be  simple  substances.  The  discovery  of  potas- 
sium by  Davy  dispelled  this  idea,  and  the  true  character  of 
these  metals  was  shown.  Davy  isolated  potassium  by  de- 
composing potash  by  means  of  a  strong  electric  current. 

Potassium  is  found 
widely  distributed  in  na- 
ture. It  forms  from  2  to 
3  per  cent  of  our  granite 
rocks.  In  combination 
with  chlorin,  as  potassium 
chlorid,  it  is  found  in  the 
earth  in  considerable  de- 
posits. No  vegetable 
growth  is  possible  with- 
out potassium ;  hence  all  fruitful  soils  contain  it. 

Manufacture.  Potassium  is  obtained  by  reducing  potas- 
sium carbonate  with  carbon.  A  mixture  of  potassium  tarta- 
rate  and  potassium  carbonate  is  first  heated.  The  tartarate 
decomposes  into  carbonate  and  carbon.  This  mixture  is 
then  placed  in  a  wrought  iron  mercury  bottle,  and  is  heated. 


A,  Grate.    B,  Retort.    C,  Receiver. 


196      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Potassium  and  carbon  monoxid  results :  K2CO3  +  2  C  =  K2 
+  3  CO.  These  two  substances  readily  form  an  explosive 
compound  at  high  temperatures,  but  this  is  now  prevented 
by  cooling  the  potassium  as  fast  as  it  is  produced.  The 
retort  is  connected,  by  means  of  a  short  tube,  with  a  very 
shallow  receiver  consisting  of  two  thin  iron  plates  clamped 
together.  The  potassium  here  condenses  to  a  liquid,  and  is 
transferred  to  a  vessel  containing  naphtha,  as  soon  as  it 
solidifies.  (See  ill.  p.  195.) 

POTASSIUM    CARBONATE 

Potassium  carbonate  was  formerly  manufactured  almost 
entirely  from  wood  ashes.  The  ashes  were  treated  with 
water,  filtered,  and  the  solution  evaporated.  Nowadays,  in 
addition  to  the  above  source,  most  of  the  "  potash  "  of  com- 
merce is  obtained  from  three  sources  :  (a)  from  beet-root, 
(£)  from  the  sweat  of  sheep,  and  (c]  from  potassium  sulfate. 
(a)  The  molasses  from  beet-root  sugar  is  allowed  to  ferment, 
and  is  then  evaporated.  A  black  mass  containing  the  potash 
is  then  obtained.  (//)  One-third  of  the  weight  of  the  sweat 
of  sheep  is  potassium  compounds.  The  washings  of  sheep 
wool  are  evaporated  to  dryness,  and  then  heated  in  retorts. 
What  is  left  is  carbon  and  various  potassium  salts,  which  are 
then  separated,  (c)  Potassium  sulfate  is  obtained  as  a  by- 
product in  many  processes.  It  is  converted  to  the  carbonate 
by  a  process  noted  later  under  the  manufacture  of  sodium 
carbonate. 

POTASSIUM    HYDROXID 

Potassium  hydroxid  is  obtained  from  the  carbonate  by 
treating  it  with  slaked  lime. 

K2  C  03  +  Ca  (0  H)2  =  2  K  0  H  +  Ca  C  03 
To  a  hot  solution   of  potassium  carbonate,   lime   is   added 


HISTORY,    OCCURRENCE    AND    APPLICATIONS 

until,  after  the  calcium  carbonate  formed  has  settled,  the 
addition  of  hydrochloric  acid  causes  no  effervescence.  The 
liquid  is  then  drawn  off,  evaporated,  and  finally  heated  to 
redness  in  silver  crucibles.  From  the  crucibles,  it  is  run 
into  cylindrical  molds,  and  thus  cast  into  sticks. 

Potassium  hydroxid,  together  with  sodium  hydroxid,  are 
used  in  the  manufacture  of  soap.  Soap  is  a  salt  of  an  alkali 
and  an  organic  acid. 

POTASSIUM    NITRATE 

This  substance  is  found  in  nature  exuding  from  the  soil  in 
warm  climates.  (See  page  179.)  It  is  also  obtained  artifi- 
cially on  what  are  called  "  niter  plantations."  A  mound  of 
chalky  soil  is  built  upon  a  foundation  of  clay.  This  is  kept 
moist  with  the  nitrogenous  refuse  from  stables  and  sewers. 
In  time  the  organic  matter  oxidizes,  and,  uniting  with  the 
alkalis,  forms  nitrates.  The  soil  is  washed  from  time  to  time 
to  obtain  these  salts.  From  the  solution  thus  obtained, 
crude  potassium  nitrate  is  separated  by  crystallization. 

Application.  The-  chief  uses  of  potassium  nitrate  are  in 
the  manufacture  of  nitric  acid  and  gunpowder. 

Gunpou>der.  Gunpowder  is  a  mixture  of  potassium  nitrate, 
charcoal,  and  sulfur.  The  percentages  are  approximately, 
potassium  nitrate  75,  charcoal  15,  and  sulfur  10.  The 
oxygen  necessary  for  the  burning  of  the  charcoal  and 
sulfur  is  furnished  by  the  potassium  nitrate.  This  combus- 
tion takes  place  very  rapidly,  forming  large  quantities  of 
carbon  dioxid  and  nitrogen ;  hence  the  explosion. 

OTHER    POTASSIUM    SALTS 

Among  other  potassium  salts  that  need  not  be  described 
here  may  be  mentioned  potassium  chlorid  (KC1,  see  Exp. 


198       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

25  c),  potassium  bromid  (KBr,  see  Exp.  28  c),  potassium 
iodid  (KI,  see  Exp.  30  «),  potassium  chlorate  (KC1O3,  see 
Exp.  6  c),  potassium  sulfid  (K2S),  and  potassium  cyanid 

(KCN). 

SODIUM 

History  and  Occurrence.  Sodium  was  first  obtained  in 
1807  by  Davy  in  the  same  manner  as  he  obtained  potassium. 
(See  page  195.)  It  never  occurs  free  in  nature,  but  in  com- 
bination it  is  very  plentiful,  occurring  in  the  sea  in  sodium 
chlorid,  and  on  the  land  chiefly  as  sodium  chlorid,  nitrate, 
carbonate,  and  sulfate. 

Manufacture  and  Uses.  It  is  prepared  in  a  manner  analo- 
gous to  that  of  potassium.  In  the  case  of  sodium,  however, 
there  is  no  liability  to  explosions,  since  sodium  does  not 
form  a  compound  with  carbon  monoxid. 

Its  chief  use  is  in  the  preparation  of  the  metals  manganese 
and  aluminum.  It  is  comparatively  cheap. 


SODIUM    CHLORID 

Occurrence  and  Extraction.  Sodium  chlorid,  as  has  been 
stated  before,  occurs  plentifully  in  the  sea  and  in  various 
salt  beds. 

It  can  be  obtained  easily  from  sea  water ;  but  by  far  the 
greater  amount  of  the  salt  of  commerce  is  either  mined  in 
the  solid  state,  as  rock  salt,  or  else  is  extracted  from  the 
earth  in  the  form  of  brine.  The  brine  is  then  evaporated. 

Uses.  Besides  its  uses  for  seasoning  food  and  preserving 
meats,  it  is  the  basis  of  the  great  soda  industry.  First  the 
sulfate  is  made  from  the  chlorid,  and  then  the  other  sodium 
compounds  are  obtained  from  the  sulfate. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         1 99 


SODIUM    SULFATE 

Occurrence.  Sodium  sulfate  occurs  in  nature  as  the  min- 
eral thenardite.  The  sulfate  of  commerce,  however,  is  made 
from  common  salt. 

Manufacture  and  Uses.  —  Sodium  chlorid  and  concentrated 
sulfuric  acid  are  mixed  in  a  large  iron  pan  or  retort,  and 
this  mixture  is  heated  in  a  reverberatory  furnace.*  The 
mixture  is  first  gently  heated,  acid  sodium  sulfate  and  hydro- 
chloric acid  being  formed. 

Na  Cl  +  H2  S  04  =  Na  H  S  04  +  H  Cl 

The  acid  gas  passes  through  a  flue  into  a  series  of  towers 
containing  coke,  through  which  water  is  trickling.  The 
water  collects  the  acid  fumes,  and  is  drawn  off  at  the  bottom, 
forming  the  hydrochloric  acid  of  commerce. 

The  mass  left  in  the  pan  is  then  transferred  to  another 

part  of  the  furnace,  and  is  subjected  to  a  higher  temperature. 

This    forms    the    normal    sodium   sulfate   (see    Defs.   under 

Exp.  1 8),  which  is  the   "  salt  cake  "  used  in  the  manufacture 

of  sodium  carbonate. 

*  A  reverberatory  furnace  is  one  that  has  two  or  more 
compartments,  one  in  which  the  fuel  is  burned,  and  others 
in  which  the  substances  to  be  treated  are  placed.     The 
compartments  are  so  arranged  that  the  flames   from    the 
burning  fuel  are  drawn  by 
the  draft  over  and  deflected 
down  upon  the  substances 
in  the  other  compartments, 
thus  producing  the  heat   to 
act  upon  them.     Fuel  that 
A,  Firebox.    B,  Hearth  upon  which  substance       gives  long  flames  is  generally 
to  be  heated  is  placed .  used. 


2OO      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


SODIUM    CARBONATE 

There  are  two  main  processes  for  the  manufacture  of 
sodium  carbonate.  The  older  process  is  called  the  Le  Blanc, 
while  the  more  recent  is  called  the  Solvay  or  the  Ammonia 
process. 

The  Le  Blanc  Process.  In  the  Le  Blanc  method,  sodium 
sulfate  is  treated  with  charcoal  (coal)  and  calcium  carbonate. 
When  heated,  the  sodium  sulfate  is  reduced  by  the  charcoal 
to  sodium  sulfid  (Na^S),  which  in  turn  acts  with  calcium 
carbonate  forming  sodium  carbonate. 

(1)  Na2  S  04  +  4  C  =  Na2  S  +  4  C  0 

(2)  Na2  S  +  Ca  C  03  =  Na2  C  03  +  Ca  S 

The  mixture  is  first  heated  in  the  coolest  part  of  a  reverbera- 
tory  furnace,  to  produce  the  first  reaction.  After  a  time  it  is 
placed  in  the  hottest  part  of  the  furnace,  when  the  second 
reaction  takes  place.  The  "  black  ash "  (a  mixture  of 
sodium  sulfid,  N^S,  sodium  carbonate,  NasCOg,  coal,  and 
lime)  thus  obtained  is  treated  with  water.  The  soluble 
sodium  carbonate  is  then  removed  from  the  solution. 

The  Solvay  Process.  For  many  years,  the  Le  Blanc  was  the 
best  process  known  for  the  manufacture  of  sodium  carbonate. 
It  has  been  demonstrated,  however,  that  the  Solvay  process 
is  more  economical. 

A  solution  of  salt  is  impregnated  with  ammonia  gas  in  the 
proper  proportions  (a  molecule  of  ammonia  for  every  mole- 
cule of  salt).  Then  into  this  there  is  led  carbon  dioxid  until 
the  solution  is  saturated.  Ammonium  chlorid  and  acid 
sodium  carbonate  are  formed  according  to  the  following 
reaction :  — 

Na  Cl  +  N  H4  H  C  03  =  N  H4  Cl  +  Na  H  C  03 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         2OI 

On  account  of  its  slight  solubility  in  the  liquid  used,  the 
acid  sodium  carbonate  separates  out  of  the  solution,  is  col- 
lected, and  dried.  The  liquid  remaining  is  treated  with 
lime,  and  the  ammonia  thus  obtained  is  used  again  for  im- 
pregnating a  new  solution  of  salt.  The  acid  carbonate  may 
be  heated,  and  the  carbon  dioxid  from  this  may  also  be  used 
again,  the  residue  left  after  heating  being  the  normal  car- 
bonate. 

Applications.  Immense  quantities  of  sodium  carbonate 
are  used  in  the  manufacture  of  glass  and  soap.  It  is  also 
used  in  the  preparation  of  other  sodium  compounds.  The 
housewife  uses  the  normal  carbonate  for  softening  water,  and 
the  acid  carbonate  *  for  cooking  purposes. 

SODIUM    HYDROXID 

After  sodium  carbonate  has  been  separated  from  the 
"  black  ash  "  obtained  in  the  Le  Blanc  process,  there  remains 
in  the  liquid  sodium  hydrate  that  has  been  formed  by  the 
action  of  the  lime  present.  The  solution  is  heated*  air  is 
blo\vn  through  it,  and  a  quantity  of  sodium  nitrate  is  added. 
This  oxidizes  the  sulfid  present  to  sulfate.  The  solution 
is  then  evaporated  to  dryness,  and  raised  almost  to  a  red 
heat.  The  cooled  product  is  commercial  sodium  hydroxid. 

Sodium  hydroxid  is  also  manufactured  by  treating  a  weak 
solution  of  sodium  carbonate  with  lime. 

*  Sodium  bicarbonate,  the  acid  salt,  is  made  by  passing  carbon 
dioxid  over  sodium  carbonate  dissolved  in  its  water  of  crystallization. 
The  principal  use  of  sodium  bicarbonate  is  in  the  manufacture  of  bak- 
ing powder.  The  constituents  of  baking  powder  are  sodium  bicarbonate 
and  some  acid,  or  acid  salt.  When  mixed  with  dough,  these  consti- 
tuents, by  the  aid  of  water  in  the  dough  and  the  heat  of  the  oven,  react 
upon  each  other.  One  of  the  products  of  the  reaction  is  carbon  dioxid  ; 
it  is  to  this  that  the  leavening  process  is  due. 


2O2      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

Na2  C  03  +  Ca  (0  H)2  =  Ca  C  03  +  2  Na  0  H 

The  calcium  carbonate,  being  insoluble,  is  removed  by 
nitration ;  and  the  filtrate,  when  evaporated,  yields  the 
hydroxid.  Sodium  hydroxid  is  purified  by  treating  the  crude 
product  with  alcohol,  which  dissolves  the  hydroxid  and 
leaves  the  impurities.  Such  sodium  hydroxid  is  called 
"  soda  by  alcohol." 

Applications.  Sodium  hydroxid  finds  its  main  uses  in  the 
soap  factory.  It  is  also  used  in  purifying  petroleum,  car- 
bolic acid,  in  the  manufacture  of  numerous  chemicals,  in  dye 
works,  and  in  the  manufacture  of  paper. 

CALCIUM 

History  and  Occurrence.  Davy,  the  discoverer  of  sodium  and 
potassium,  first  prepared  calcium  as  a  powder  by  electrolysis  ; 
but  Matthiessen  obtained  the  first  "  piece  of  calcium "  in 
1856.  The  metal  itself  does  not  occur  free  in  nature,  nor  is 
it  of  any  importance.  Its  various  compounds,  however,  are 
very  plentiful,  forming,  as  dolomite,  CaMg(CO3)2,  whole 
mountain  ranges.  It  is  also  found  in  large  quantities  as  the 
sulfate  and  phosphate.  The  bones  of  animals  and  the 
shells  of  eggs  and  of  mollusks  are  mainly  composed  of  cal- 
cium salts.  In  fine,  it  is  one  of  the  most  plentiful  of  the 
elements. 

CALCIUM    CARBONATE 

Occurrence  and  Application.  This  compound  occurs  as 
Iceland  spar,  in  transparent  crystals.  The  Carrara  marble 
is  a  very  pure  form  of  calcium  carbonate,  while  ordinary 
limestone  contains  various  impurities.  The  artificial  sub- 
stance is  made  by  treating  the  chlorid  with  ammonium 
carbonate. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         2O3 

The  mineral  is  used  for  building  purposes,  and  for  the 
manufacture  of  lime  and  cements.  The  beautiful  marbles  of 
Italy  are  used  mostly  by  the  sculptor. 

CALCIUM    OXID 

Manufacture.  The  natural  carbonate  is  heated  in  special 
furnaces,  called  kilns.  Fuel  is  allowed  to  burn  under  the 
limestone,  thus  driving  off  the  carbon  dioxid  and  leaving  the 
calcium  oxid  (see  Exp.  21,  Part  I.). 

Applications.  The  chief  use  of  lime  is  in  the  making  of 
mortar  and  cements. 

Common  mortar  is  a  mixture  of  one  part  of  slaked  lime 
and  three  or  four  parts  of  sand  made  into  a  pasty  mass. 
This  hardens  or  "  sets  up  "  in  a  few  days,  but  it  takes  years 
for  it  to  harden  completely.  This  is  first  due  to  the  evapora- 
tion of  the  water ;  then,  as  time  goes  on,  the  lime  takes  on 
carbon  dioxid  from  the  air,  and  becomes  calcium  carbonate. 

Hydraulic  mortar  is  made  from  lime  containing  more  than 
10  per  cent  of  silica,  and  has  the  ability  to  harden  under 
water. 

Portland  cement  is  a  hydraulic  mortar  made  from  chalk 
and  clay.  The  two  are  ground  together  in  water,  dried,  and 
burnt  in  kilns. 

CALCIUM    HYDROXID 

Calcium  hydroxid  is  simply  slaked  lime  (see  Exp.  20  £, 
Part  I.). 

CALCIUM    SULFATE 

Occurrence  and  Application.  Calcium  sulfate  occurs  as 
gypsum  with  water  of  crystallization,  and  as  anhydrite 
without  it.  The  artificial  salt  is  formed  by  the  action  of 
sulfuric  acid  on  the  carbonate. 


2O4        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

After  gypsum  has  been  heated  so  as  to  lose  its  water  of 
crystallization,  it  is  called  "  plaster  of  Paris,"  and  possesses 
the  property  of  hardening  when  moistened  with  water.  It 
is  therefore  extensively  used  for  making  plaster  casts,  and 
as  a  cement.  Gypsum  is  found  also  in  some  fertilizers. 

CALCIUM    CHLORID 

Calcium  chlorid  is  obtained  as  a  by-product  in  many 
manufacturing  processes.  On  account  of  its  strong  attrac- 
tion for  water,  it  is  used  in  laboratories  for  the  drying  of 
gases. 

BARIUM    AND    STRONTIUM 

History  and  Occurrence.  The  elements  barium  and 
strontium  were  isolated  by  Davy  in  1806,  in  connection 
with  his  experiments  on  calcium.  They  occur  chiefly  as 
barite  (BaSO4),  witherite  (BaCO3),  strontianite  (SrCO3),  and 
celestite  (SrSO4). 

Compounds.     See  Exps.  58  and  59,  Part  I. 

MAGNESIUM 

History  and  Occurrence.  Magnesium  also  was  discovered 
by  Davy.  It  occurs  as  dolomite,  MgCa(CO3)2,  on  the  earth ; 
as  the  sulfate,  MgSO4,  in  mineral  springs  ;  and  as  chlorid, 
MgCl2,  in  the  sea.  The  metal  is  obtained  by  reducing  the 
chlorid  by  means  of  sodium. 

Compounds.      See  Exps.  13  and  17  b,  Part  I. 

Applications.  Magnesium  is  used  in  many  chemical  oper- 
ations, and  as  a  means  of  artificial  light  in  photography. 

The  sulfate  (Epsom  salts)  and  the  carbonate  are  exten- 
sively used  in  medicine.  The  chlorid  is  used  in  the  manu- 
facture of  cotton  goods. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         20$ 


ZINC 

History  and  Occurrence.  The  alloy  brass  was  known  to 
the  ancients,  but  they  did  not  know  that  it  contained  any 
metal  besides  copper.  The  discoverer  of  the  second  metal, 
zinc,  is  uncertain,  although  it  is  mentioned  in  writings  of 
the  1 6th  century.  It  is  found  plenteously  as  the  sulfid 
(ZnS,  zinc  blende),  as  the  oxid  (ZnO,  red  zinc  ore),  as  the 
carbonate  (ZnCO3,  calamine),  and  as  the  silicate  (  (ZnO)2SiO2, 
H2O,  electric  calamine). 

Extraction.  The  ore  is  first  roasted  to  convert  it  into  oxid. 
The  oxid  is  then  mixed  with  carbon.  It  is  then  placed  in 
cylindrical  fire-clay  retorts  about  three  feet  long  and  eight 
inches  in  diameter.  These  retorts  are  arranged  in  tiers  and 
set  slantingly.  To  the  open  end  of  each  retort  is  joined  a 
conical  receiver  about  10  inches  long,  extending  downwards. 
The  retorts  are  heated,  and  soon  burning  carbon  monoxid 
appears  at  the  opening  of  each  receiver.  The  characteristic 
greenish-blue  flame  soon  appears,  showing  that  the  metal 
is  volatilizing.  The  reduced  metal  is  removed  from  time  to 
time,  the  whole  operation  requiring  about  eleven  hours. 
The  zinc  is  afterwards  redistilled  to  purify  it. 

Industrial  Applications.  Zinc  is  largely  used  in  "  galvaniz- 
ing "  iron,  which  is  done  by  dipping  clean  iron  into  melting 
zinc.  It  is  also  used  in  the  manufacture  of  brass  and  other 
alloys.  Zinc  dust  finds  an  extensive  use  in  organic  chem- 
istry. It  is  used  on  a  large  scale  in  the  manufacture  of 
indigo  blue.  It  is  also  used  as  a  paint  for  iron  articles. 

Compounds.     See  Exps.  12  and  15  £,  Part  I. 

Uses.  The  oxid  (ZnO)  is  largely  used  as  a  white  paint. 
The  chlorid  (ZnCL)  is  a  strong  antiseptic,  and  is  used  also 
in  soldering.  The  sulfate  (ZnSO4)  is  used  in  dyeing  and 
calico  printing. 


2O6       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


CADMIUM 

History  and  Occurrence.  Cadmium  was  discovered  by 
Stromeyer  in  1817.  It  occurs  in  zinc  ores.  The  first  batch 
of  zinc  obtained  from  the  zinc  smelter  mentioned  above  con- 
tains practically  all  the  cadmium  ;  and,  since  it  boils  at  a 
lower  temperature  than  zinc,  the  cadmium  distills  over  first. 

Compounds.     See  Exp.  51,  Part  I. 


ALUMINUM 

History  and  Occurrence.  Aluminum  was  first  obtained  by 
Wohler  in  1827,  who  separated  the  metal  by  means  of 
metallic  sodium.  It  is  now  obtained  by  an  electrolytic 
method  invented  by  Hall  in  1886. 

Aluminum  occurs  almost  entirely  as  a  silicate.  Clay, 
mica,  slate,  etc.,  are  all  silicates  of  aluminum.  Its  oxid 
(A12O3)  is  a  common  ore,  and  cryolite  (AlNa3F6)  is  a  very 
useful  mineral.  The  precious  stones,  ruby  and  sapphire, 
are  oxids  of  aluminum,  colored  respectively  by  chromium 
and  cobalt  compounds. 

Extraction.  Hall's  method  of  reducing  aluminum  com- 
pounds is  as  follows  :  A  large  iron  receptacle  is  lined  with 
carbon.  Into  this  extend  a  number  of  large  carbon  elec- 
trodes. A  mixture  consisting  of  cryolite  and  fluorite  is 
placed  in  the  receptacle.  A  strong  electric  current  is  then 
made  to  pass  through  the  apparatus.  The  current  passes 
in  through  the  carbon  rods,  which  act  as  the  positive  elec- 
trode, while  the  carbon  lining  acts  as  the  negative  electrode. 
After  the  flux  is  melted  by  the  current,  aluminum  oxid 
(A12O8)  is  added  at  intervals.  The  melted  aluminum  collects 
at  the  negative  electrode,  while  the  oxygen  unites  with 


HISTORY,    OCCURRENCE   AND    APPLICATIONS         2O/ 

carbon  at  the  positive.  The  aluminum,  when  it  has  formed 
in  sufficient  quantities,  is  ladled  out  from  the  flux. 

Applications.  On  account  of  its  lightness,  aluminum  is 
used  for  the  manufacture  of  utensils  where  great  strength  is 
not  required.  It  forms  numerous  alloys,  one  of  which,  alumi- 
num bronze,  is  very  beautiful. 

Compounds  of  Aluminum.     See  Exp.  55,  Part  I. 

Manufacture  of  Aluminum  Sulfate.  Finely  powdered 
China  clay,  aluminum  silicate,  is  roasted  and  then  heated 
with  sulfuric  acid.  In  case  other  clays  are  used,  the  iron 
is  precipitated  by  means  of  potassium  ferrocyanid.  The 
product  is  sold  on  the  market  under  the  name  of  "  alum 
cake." 

Alum.  An  alum  is  a  double  sulfate  of  a  triad  metal 
and  an  alkali,  and  has  twelve  molecules  of  water  of  crys- 
tallization. Aluminum  sulfate  is  not  readily  obtained  in 
crystalline  form ;  hence  it  is  usually  crystallized  together 
with  potassium  or  ammonium  sulfate,  with  which  it  forms 
crystals  of  potassium  or  ammonium  aluminum  sulfate,  KA1 
(SO4)2,i2H2O,  in  the  form  of  regular  octahedrons. 

The  best  commercial  method  for  obtaining  alum  is  to 
roast,  together  with  coal,  the  bituminous  shale  of  coal  beds. 
This  shale  is  generally  rich  in  aluminum  silicate  and  iron 
pyrites.  The  ferric  sulfid  gives  off  sulfur,  which  becomes 
oxidized  to  sulfuric  acid.  Both  aluminum  and  ferric  sul- 
fates  are  formed.  This  mass  is  dissolved  in  water,  and 
evaporated  down  until  the  ferric  sulfate  separates  out. 
Then  crude  potassium  chlorid  and  potassium  sulfate  are 
added,  and  the  mixture  is  agitated  until  it  is  cold.  The 
small  crystals  then  formed  are  washed  with  cold  water,  and 
recrystallized. 

Uses  of  Alum.  Alum  is  used  mostly  in  dyeing  establish- 
ments as  a  mordant.  The  cloth  is  placed  in  an  alum  solution 


2O8        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

to  which  sodium  carbonate  has  been  added  until  the  pre- 
cipitate first  formed  has  been  redissolved.  After  the  cloth 
is  dry  the  color  is  applied.  The  mordant  causes  the  color 
to  stick  to  the  fibers  of  the  material. 

Alum  is  also  used  in  the  manufacture  of  paper. 

CHROMIUM 

History  and  Occurrence.  Vanquelin  and  Klaproth  dis- 
covered, independently  of  each  other,  the  element  chromium 
in  the  mineral  crocoisite.  It  was  named  chromium  (from 
Xpw/oa,  color),  because  its  compounds  are  all  colored. 

It  is  found  in  crocoisite  (PbCrO4),  and  in  chromite  (Fe 
O,Cr2O3).  The  green  color  of  emeralds  is  due  to  the  presence 
of  chromium. 

It  is  a  hard,  gray,  almost  infusible  metal. 

Compounds.  Its  principal  compounds  are  chrome  alum, 
CrK(SO4)2,i2H2O,  used  in  dyeing  and  tanning;  lead  chro- 
mate,  PbCrO4,  a  valuable  pigment;  potassium  chromate, 
K2CrO4;  and  potassium  bi-chromate,  K2Cr2O7.  The  last 
is  the  most  important  commercial  product,  and  is  used  ex- 
tensively in  bi-chromate  batteries  and  in  dyeing. 

IRON 

History.  It  is  supposed  that  iron,  the  most  important  and 
most  useful  of  all  the  metals,  was  first  extracted  from  its 
ores  in  India.  Moses  speaks  of  iron  as  used  by  the  Hebrews. 
The  Greeks  obtained  their  iron  from  the  Chalybes,  a  nation 
living  on  the  shores  of  the  Black  Sea.  The  Romans  obtained 
theirs  from  Spain  and  Elba,  and  from  their  own  mines.  Our 
own  iron  is  mostly  obtained  from  Pennsylvania,  although 
great  iron  industries  have  sprung  up  in  Alabama,  Ten- 
nessee, and  Illinois. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         2OQ 

Occurrence.  Iron  is  found  widely  distributed  in  the  earth. 
It  occurs  free  in  certain  rocks  and  in  meteorites.  The 
common  ores  of  iron  are:  —  magnetite  (Fe3O4),  haematite 
(Fe2O3),  siderite  (FeCO3),  and  limonite 


A,  Tuyere.    B,  Hearth.    C,  Molten  iron.    D,  Limestone,  coke,  and  ore. 
E,  Sand. 

Extraction.  It  would  seem  to  be  a  very  simple  process  to 
obtain  iron  from  its  ore.  Theoretically,  all  that  is  necessary 
with  most  ores  would  be  to  remove  the  oxygen.  This  is 
done  in  practice  by  means  of  carbon,  but  difficulties  come  up 
that  have  to  be  conquered  before  the  final  product  is  ob- 
tained. Theoretically,  carbon  would  be  the  only  thing 
necessary  to  use  to  remove  the  oxygen,  thus :  — 

Fe3  04  +  2  C  =  3  Fe  +  2  C  02 


2IO        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

It  is  found  that  carbon  monoxid  is  also  formed  thus :  — 

2  Fe3  04  +  6  C  =  6  Fe  +  4  C  0  +  2  C  02 

But  carbon  monoxid  is  also  a  reducing  agent  when  the  tem- 
perature is  high  enough,  and  acts  upon  the  ore,  thus :  — 

Fe3  04  +  4  C  0  =  3  Fe  +  4  C  02 

Iron  ore,  however,  is  rarely  pure,  the  impurities  being 
mostly  silicates  in  some  form  or  other.  These  must  be 
fused  or  the  carbon  cannot  get  a  chance  to  act  upon  the  ore. 
To  effect  this,  limestone  (CaCO8)  is  added  to  the  carbon, 
which  is  used  in  the  form  of  coke.  The  limestone,  when 
heated  together  with  the  silicates,  produces  a  flux  or  molten 
mass.  The  process  is  carried  on  in  a  furnace  from  75  to 
100  ft.  high.  The  ore,  coke,  and  limestone  are  shoveled 
into  the  furnace  in  layers  and  ignited.  A  blast  of  air  is 
driven  through  the  furnace  from  the  bottom  by  means  of  a 
number  of  pipes  called  tuyeres.  The  reduction  takes  place ; 
the  iron,  from  its  greater  specific  gravity,  sinks  to  the  bottom 
of  the  furnace,  and  is  drawn  off  into  molds  of  sand  three  or 
four  feet  long  and  three  or  four  inches  wide  and  deep.  This 
is  called  "pig  iron,"  and  contains  as  impurities  sulfur,  sili- 
con, carbon,  manganese,  and  a  small  amount  of  phosphorus. 

Wrought  Iron.  The  pig  iron  obtained  from  the  blast  fur- 
nace is  made  into  wrought  iron  by  removing  from  it  the 
carbon,  of  which  there  is  from  2  to  6  per  cent.  The  process 
is  called  "  puddling."  The  pig  iron  is  placed  in  a  reverbera- 
tory  furnace.  The  sides  of  the  furnace  are  lined  with  a  sub- 
stance containing  haematite.  At  first  the  iron  is  melted 
slowly.  Finally  it  boils,  and  some  of  the  carbon  is  removed 
by  the  oxygen  of  the  haematite.  The  heat  is  then  in- 
creased, and  the  molten  mass  is  stirred  by  workmen  by 
means  of  iron  rods.  When  the  carbon  is  burned  away,  the 
iron  becomes  pasty,  and  is  removed  in  large  masses  on  the 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         21  I 


ends  of  the  rods.  The  iron  thus  obtained  is  tough.  Most 
of  the  phosphorus,  sulfur,  and  silicon  of  the  pig  iron  remains 
in  the  furnace  in  the  slag,  and  is  afterwards  removed. 

Steel.  Steel  can  be  made  from  either  pig  iron  or  wrought 
iron. 

To  make  it  from  wrought  iron,  the  bars  are  packed  in 
powdered  charcoal,  and  heated  for  a  number  of  days  at  a 
red  heat.  The  bars  are  then  allowed  to  cool  slowly.  The 
best  quality  of  steel  is  made  in  this  way.  It  is  called 
"  blister  steel,"  since,  when 
taken  from  the  furnace,  the 
bars  are  covered  with  blis- 
ters. 

The  Bessemer  process  is 
the  ordinary  way  for  making 
steel.  Good  pig  iron  is 
melted  in  huge  vessels, 
called  converters,  hung  on 
trunnions.  A  blast  of  air 
is  blown  through  the  molten 
iron  from  the  bottom.  The 
flame  is  watched  by  an 
expert  until  just  the  right  moment,  and  then  the  blast  is 
stopped.  Carbon  is  then  added  to  the  purified  iron  by 
means  of  molten  spiegeleisen,  a  kind  of  cast  iron  that  contains 
carbon  and  manganese.  The  converter  is  then  emptied,  and 
the  steel  is  cast  into  ingots. 

PROPERTIES  OF  THE  VARIOUS  KINDS  OF  IRON 

Pig  iron  is  brittle  and  cannot  be  welded.  It  is  the  kind 
of  iron  used  in  making  castings.  Wrought  iron  is  tough, 
malleable,  ductile,  and  can  be  welded.  Steel  is  usually  brittle. 
Its  hardness  can  be  changed  by  tempering;  that  is,  by 


212        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

heating  at  various  temperatures  and  suddenly  cooling.  By 
allowing  hot  steel  to  cool  slowly,  it  can  be  made  almost 
as  soft  as  wrought  iron. 

Compounds  of  Iron.  See  Exps.  14  a,  b,  15  /  and  Exp. 
56,  Part  I. 

FERROUS    SULFATE 

One  source  of  ferrous  sulfate  is  the  bituminous  shale  from 
coal  beds.  The  ferrous  sulfate  is  crystallized  out  from  the 
solution  of  the  roasted  product.  (See  manufacture  of  alum.) 
It  is  used  in  the  manufacture  of  black  ink,  in  certain  photo- 
graphic processes,  and  as  a  reagent  in  the  laboratory. 

POTASSIUM    FERROCYANID    AND    FERRICYANID 

When  nitrogenous  organic  matter  is  fused  with  potassium 
hydroxid  and  iron  filings,  and  the  product  is  treated  with 
water,  a  solution  is  obtained  which,  when  purified  and  evapo- 
rated, gives  yellow  crystals  of  potassium  ferrocyanid,  K4Fe 
(CN)6.  This  is  commonly  called  yellow  prussiate  of  potash. 
If  treated  with  ferric  chlorid  a  beautiful  blue  precipitate, 
soluble  in  excess  of  the  chlorid,  is  formed  and  is  called 
Prussian  blue. 

If  the  potassium  ferrocyanid  is  treated  with  chlorin,  the 
ferricyanid  is  formed  according  to  the  formula :  — 

2  K4  Fe  (C  N)6  +  C12  =  2  K  Cl  +  2  K3  Fe  (C  N)6 
When  evaporated  and  crystallized,  red  crystals  of  the  ferri- 
cyanid are  obtained.     No  precipitate  with  ferric  compounds 
is  formed  by  the  ferricyanid,  but  with  ferrous  salts   a  dark 
blue  precipitate  called  TurnbulVs  blue  is  formed. 

MANGANESE 

History  and  Occurrence.  In  1807,  Gahn  first  isolated 
manganese.  Its  most  abundant  and  best  known  ore  is  pyro- 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         213 

lusite    (MnO2).     The  metal  itself  is  reddish-white,  oxidizes 
easily,  and  must  be  kept  away  from  contact  with  air. 

Applications.  It  finds  its  main  use  in  its  compound  man- 
ganese dioxid  (MnO2),  which  is  a  \^ery  valuable  oxidizing 
agent.  Spiegeleisen  is  an  alloy  with  iron  and  contains  car- 
bon. As  has  been  stated  before,  spiegeleisen  is  used  in  the 
manufacture  of  steel  from  pig  iron.  Potassium  permanganate 
is  also  a  valuable  oxidizing  agent.  Manganese  sulfate  is 
often  used  as  a  mordant. 

NICKEL   AND    COBALT 

History.  Nickel  was  discovered  by  Cronstedt  in  1751, 
and  cobalt  probably  by  Brandt  somewhere  about  1735. 
The  two  metals  both  occur  in  combination  with  arsenic  as 
arsenids.  They  are  almost  always  associated,  and  are  sepa- 
rated with  difficulty.  Of  the  two  metals,  nickel  is  the  more 
important. 

Extraction  of  Nickel.  The  ore,  nickel  arsenid,  is  heated  in 
a  reverberatory  furnace,  thus  driving  off  most  of  the  arsenic. 
The  residue  is  dissolved  in  hydrochloric  acid ;  and  ftie 
impurities,  with  the  exception  of  cobalt,  are  removed  by 
various  processes.  From  the  remaining  solution,  the  cobalt 
is  precipitated  as  Co(OH)2  by  bleaching  powder.  When 
heated,  the  cobalt  hydroxid  becomes  cobalt  oxid,  and  is 
removed  by  filtering.  The  nickel  solution  is  then  treated 
with  calcium  hydrate,  and  is  finally  reduced,  at  a  white  heat, 
to  nickel  by  means  of  carbon.  It  appears  on  the  market  in 
small  cubes. 

Uses.  Nickel  is  valuable  as  an  alloy  with  steel,  and  with 
copper  and  zinc  it  forms  the  well-known  alloy  German  silver. 
It  is  also  used  for  plating  iron,  for  coins,  and  for  laboratory 
utensils. 

Principal  Compounds.     See  Exp.  57,  Part  I. 


214      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

TIN 

History  and  Occurrence.  Tin  was  known  to  the  Romans, 
who  obtained  it  from  England  after  Caesar's  conquest.  It 
occurs  chiefly  as  cassiterite,  or  tin  stone,  from  which  the  tin 
of  commerce  is  obtained.  The  oldest  tin  mines  are  those  of 
Cornwall.  Tin  is  also  found  in  considerable  quantities  in 
Australia  and  in  the  Black  Hills  of  South  Dakota. 

Extraction.  The  tin  ore  is  first  crushed.  It  is  then 
washed,  and  the  cleansed  ore  is  roasted,  to  free  it  from 
sulfur  and  arsenic.  Mixed  with  coal  and  a  small  quantity 
of  lime,  it  is  then  heated.  Metallic  tin  separates  out.  The 
metal  obtained  is  somewhat  impure;  and  in  order  to  remove 
the  other  metals,  it  is  heated  gently.  The  tin  that  melts 
first  is  ladled  out,  and  this  product  is  further  purified  by 
being  melted  and  then  stirred  with  wet  wooden  sticks.  The 
impurities  separate  out  and  the  metal  is  cast  into  ingots. 

Uses.  The  most  important  application  of  tin  is  in  the 
manufacture  of  tin  plate.  Its  important  alloys  are  bronze 
(tin  and  copper),  and  plumber's  solder  (tin  and  lead).  It  is 
also  rolled  into  thin  sheets  known  as  tin-foil. 

Compounds.     See  Exp.  54,  Part  I. 

Reducing  Action.  Stannous  chlorid  is  a  very  valuable 
reducing  agent  in  the  laboratory. 

BISMUTH 

History  and  Occurrence.  Bismuth  was  spoken  of  by  Basil 
Valentine  in  the  i5th  century.  Pott  first  made  a  careful 
study  of  its  properties  in  1739. 

It  occurs  in  nature  and  often  nearly  pure.  When  impure, 
it  is  mixed  with  iron,  carbon,  and  slag,  and  is  then  heated 
in  pots.  The  purer  bismuth  settles  to  the  bottom.  It  may 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         215 

be  further  purified  by  melting  it  on  an  inclined  iron  plate, 
when  the  pure  bismuth  melts  first  and  runs  off. 

Applications.  Bismuth  is  added  to  certain  alloys,  to  lower 
their  melting  point.  It  is  valuable  for  making  fuses  that 
melt  at  a  moderate  elevation  of  temperature,  as  in  automatic 
fire  extinguishers  ;  also  for  safety  fuses  and  lightning  ar- 
resters in  connection  with  the  use  of  the  electric  current. 

Compounds.     See  Exp.  50,  Part  I. 

Basic  Bismuth  Nitrate  is  extensively  used  in  medicine  in 
the  treatment  of  cholera.  It  is  also  used  as  a  cosmetic,  and 
in  the  manufacture  of  porcelain. 

LEAD 

History  and  Occurrence.  Lead  is  first  mentioned  in  the 
Book  of  Job,  and  Pliny  pointed  out  the  distinction  between 
lead  and  tin.  Lead  is  rarely  found  free  in  nature,  but  occurs 
in  large  quantities  as  galena  (PbS)  in  England,  Spain,  and 
the  United  States. 

Reduction  from  the  Ore.  The  ore  galena  is  mixed  wfth 
lime,  and  is  at  first  heated  gently,  a  current  of  air  being 
drawn  through  the  furnace.  It  is  then  heated  to  a  higher 
temperature,  and  metallic  lead  is  drawn  off.  The  reactions 
that  occur  are  :  — 

Some  lead  sulfid  is  changed  to  lead  oxid, 

2  Pb  S   +  3  02  =  2  Pb  0  +-  2  S  02 

and  some  to  lead  sulfate. 

Pb  S  +  2  02  =  Pb  S  0* 

Lead  oxid  and  lead  sulfate  each  react  upon  lead  sulfid, 
forming  lead  and  sulfur  dioxid. 

Pb  S  +  2  Pb  0  ^  3  Pb  +  S  02 
and 

Pb  S  +  Pb  S  04  =  2  Pb  +  2  S  02 


2l6      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

There  is  probably  always  some  silver  in  galena.  This,  if 
sufficient  in  amount,  is  separated  by  a  process  described 
later  (see  page  219). 

Applications.  Lead  is  used  principally  for  making  water 
pipes,  roofing,  shot,  and  various  alloys,  such  as  type-metal, 
solder,  and  pewter.  In  making  shot,  some  arsenic  is  added 
to  the  lead.  Lead  compounds  are  used  in  making  white  lead, 
and  in  glass  making. 

Compounds.  Besides  the  compounds  studied  in  Exp.  53, 
Part  I.,  there  are  a  few  others  that  should  be  noted.  These 
are  red  lead  (Pb3O4) ;  white  lead,  which  is  a  basic  lead  car- 
bonate ;  and  lead  acetate,  commonly  called  "  sugar  of  lead." 

Red  Lead  is  made  by  heating  litharge  (PbO)  on  trays  in 
a  reverberatory  furnace. 

White  Lead  is  made  by  placing  rolls  of  perforated  sheet 
lead  in  earthenware  pots  partly  filled  with  diluted  vinegar 
(acetic  acid).  The  lead  is  not  in  contact  with  the  vinegar. 
The  pots  are  covered  with  tan  bark  or  manure.  In  time  the 
acid  fumes  change  the  lead  to  lead  acetate,  and  the  carbon 
dioxid  from  the  tan  bark  or  manure  changes  this  to  the 
carbonate. 

Lead  Acetate  is  made  by  dissolving  litharge  in  acetic  acid. 

COPPER 

History  and  Occurrence.  Copper  was  probably  the  first 
metal  used  by  mankind,  having  been  used  by  pre-historic 
man  in  making  his  weapons.  It  is  very  abundant,  which  is 
fortunate,  since  it  is  one  of  the  most  useful  metals  we  have. 
It  occurs  native  in  the  Lake  Superior  regions,  the  mines 
there  furnishing  almost  chemically  pure  copper.  In  Mon- 
tana, Idaho,  and  Arizona,  vast  quantities  of  copper  occur, 
associated  with  the  precious  metals.  It  is  also  found  as  the 
oxid  in  Russia  and  Australia,  and  as  the  sulfid  in  England. 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         21  j 

Extraction.  When  the  oxid  ore  is  heated  with  carbon, 
the  metal  separates  out.  The  sulfid  ore,  however,  must  first 
be  roasted  in  order  that  part  of  the  sulfid  may  be  changed 
to  oxid.  Then  the  mixture  of  sulfid  and  oxid  is  heated 
further  in  a  reverberatory  furnace,  where  the  following  re- 
action occurs. 

Cu  S  +  2  Cu  0  =  3  Cu  +  S  02. 

Industrial  Applications.  Since  the  development  of  elec- 
tricity, copper  has  become  one  of  the  staple  products,  on 
account  of  its  use  as  an  electric  conductor.  Its  use  in  the 
preparation  of  the  alloys,  brass  and  bronze,  has  already  been 
mentioned.  As  sheathing  for  ships,  etc.,  large  quantities  are 
used.  Copper  coins  are  familiar  to  us  all. 

Compounds.     See  Exps.  4;  9  g ;  39  ;  Part  I. 

Copper  sulfate  is  an  important  commercial  product  The 
sulfid  is  gently  roasted,  and  the  copper  sulfate  thus  formed 
is  dissolved  out  with  water.  It  is  also  formed  as  a  by-pro- 
duct in  the  purifying  of  gold  and  silver.  It  is  used  in  the 
preparation  of  copper  arsenite  (Paris  green),  in  the  manu- 
facture of  other  pigments,  and  also  in  calico  printing. 

MERCURY 

History.  Mercury  is  mentioned  by  the  later  Greek  writers 
under  the  name  v&pdyvpos  (from  vS<op,  water,  and  apyvpos, 
silver).  Pliny  called  it  hydrargyrum.  The  alchemists  be- 
lieved that  mercury  was  a  component  of  all  metals.  Braune 
was  the  first  to  recognize  it  as  a  true  metal.  In  the  winter 
of  1759,  he  solidified  it  by  using  snow  and  nitric  acid  to  pro- 
duce the  necessary  cold. 

Occurrence  and  Extraction.  Mercury  occurs  free  in  small 
quantities  but  is  found  chiefly  in  combination  with  sulfur  as 
cinnabar  (HgS),  in  Idria,  Carrolia,  in  Spain,  and  in  Cali- 
fornia. 


2l8        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

The  ore  mixed  with  lime  is  placed  in  a  furnace  into  which 
a  current  of  air  may  pass.  When  the  ore  is  heated,  the  sul- 
fur becomes  sulfur  dioxid ;  this,  together  with  the  vapor  of 
mercury,  passes  into  a  series  of  chambers.  The  mercury 
vapor  is  here  condensed,  while  the  sulfur  dioxid  passes 


A,  Firebox.    B,  Ore.    0,  Condensing  chambers.    D,  Chambers  through 
which  air  is  admitted. 

on.  In  the  last  chamber  falling  water  condenses  the  last 
portion  of  the  vapor.  The  mercury  thus  obtained  is  par- 
tially purified  by  being  filtered  through  linen,  and  then  by 
the  action  of  dilute  nitric  acid.  It  is  then  further  purified 
by  redistillation,  and  by  being  forced  through  chamois  skin. 
The  metal  is  usually  stored  in  wrought  iron  bottles. 

Compounds.     See  Exp.  52,  Part  I. 

Uses.  In  mining,  mercury  is  used  in  extracting  gold  and 
silver  from  their  native  mixtures.  In  the  arts,  it  is  used  in 
making  barometers,  thermometers,  mirrors,  etc.  Its  com- 
pounds are  also  used  in  the  arts  and  in  medicine.  The 
chemist  and  physicist  would  be  at  a  great  disadvantage 
without  this  convenient  metal. 

Amalgams.  Mercury  forms  amalgams  with  all  well-known 
metals  except  iron.  The  most  important  are  with  potas- 
sium, sodium,  cadmium,  copper,  silver,  and  gold.  It  also 
forms  an  amalgam  with  the  metalloid  ammonium  (see 
Exp.  42,  Part  I.). 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         2 19 

SILVER 

History  and  Occurrence.  Silver  was  one  of  the  seven 
metals  of  the  ancients.  It  occurs  in  nature,  masses  of  it 
having  been  found  weighing  as  much  as  800  Ibs.  In  a 
Norwegian  museum  is  a  mass  weighing  500  Ibs.  Its  most 
important  ores  are  argentite  (Ag.2S),  pyrargcntite  (Ag3SbS3) 
and  horn  silver  (AgCl).  Galena  generally  contains  small 
quantities  of  silver  sulfid.  The  best  known  silver  mines  are 
in  the  western  part  of  the  United  States,  in  Mexico,  in 
South  America,  and  in  Australia. 

Extraction.  The  principal  methods  of  obtaining  silver 
from  its  ores  are  the  amalgamation  process  and  cupellation 
process. 

In  the  amalgamation  process,  the  ore  is  first  roasted  with 
common  salt,  giving  silver  chlorid.  This  is  then  agitated 
with  scrap  iron  and  mercury,  giving  ferrous  chlorid  and  an 
amalgam.  The  amalgam  is  washed  and  filtered  through 
canvas,  to  remove  the  excess  of  mercury.  The  pasty  amal- 
gam that  is  left  is  then  distilled,  leaving  the  pure  silveij 
behind. 

The  cupellation  process  is  adapted  to  lead  ores  contain- 
ing silver.     These  are  first  treated 
as  described  under  lead  (see  page 
216).       In     connection     with     this 
process,    the    fol- 
lowing       prelimi- 
nary treatment  is 
necessary.        The 
lead    obtained    is 
melted    and  al- 
lowed    to     cool. 

The     verv      rmrr         A,  Fire  box.    B,  Hearth.    C,  Tuyeres.    D,  Hood. 
E,  Opening  for  introducing  materials. 


22O       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

lead  crystallizes  at  the  top,  and  is  skimmed  off.  New  alloy 
is  added,  melted,  and  the  same  operation  is  continued. 
This  is  repeated  until  fairly  pure  silver  remains.  The  silver 
thus  obtained  is  then  placed  in  a  shallow  furnace,  called  a 
cupel,  made  of  a  mixture  of  clay  and  lime.  The  alloy  is 
subjected  to  a  high  temperature,  while  a  current  of  air  is 
driven  over  the  molten  mass.  This  oxidizes  the  lead  to 
litharge,  and  pure  silver  remains. 

Applications.  Silver  has  been  used  for  coinage  from 
earliest  times.  This,  and  the  manufacture  of  silver  table- 
ware, ornaments,  etc.,  consume  most  of  the  silver  of  the 
world.  A  large  amount  of  silver  salts  is  used  in  photog- 
raphy. 

Photography.  The  fact  that  light  blackens  many  silver 
compounds  has  led  to  the  foundation  and  development  of 
the  art  of  photography.  The  first  "  light  pictures "  were 
made  by  Wedgewood  in  1802.  These  were  prints  of  leaves 
on  paper  moistened  with  silver  nitrate.  Exposure  to  light, 
however,  soon  blackened  and  destroyed  these  prints.  In 
1826,  Daguerre  found  a  process  by  which  the  image  could 
be  preserved.  Our  old  family  Daguerreotypes  are  the  sur- 
viving relics  of  this  process. 

The  present  dry  plate  method  makes  use  of  two  processes  : 
first,  the  making  of  the  negative  upon  a  sensitive  film; 
second,  the  transferring  of  the  picture  to  a  second  sensitive 
film  on  paper,  called  the  positive. 

The  negative  is  a  film  usually  composed  of  gelatine  and 
silver  bromid  (AgBr)  with  traces  of  other  silver  salts,  ad- 
hering to  a  plate  of  glass  or  some  other  transparent  sub- 
stance. When  placed  in  the  camera  and  exposed  to  light, 
no  picture  is  shown  until  the  plate  is  " developed."  This  is 
done  by  various  methods.  One  good  developing  solution 
consists  of  pyrogallic  acid  and  sodium  carbonate.  After 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         221 

the  picture  has  been  brought  out  by  the  "  developer,"  the 
unchanged  silver  bromid  must  be  removed.  This  is  done 
by  means  of  a  solution  of  sodium  thiosulfate  (NajSgOg). 
The  products  of  the  chemical  changes  are  washed  away 
with  water,  and  there  is  left  only  silver  on  the  plate.  The 
varying  amount  of  silver  on  different  parts  of  the  plate  pro- 
duces the  dark  and  light  portions. 

The  picture  must  now  be  transferred  to  paper.  One  kind 
of  sensitive  paper  has  on  one  side  of  it  a  film  of  albumen 
and  silver  chlorid.  The  negative  is  placed  over  the  paper, 
thus  leaving  the  sensitive  silver  chlorid  unprotected  from  the 
light  where  the  silver  is  least  deposited  on  the  negative.  By 
allowing  this  to  stand  in  sunlight,  the  picture  is  "  printed." 
It  must  then  be  "  toned  "  and  "  fixed."  In  "  toning,"  part  of 
the  silver  chlorid  is  usually  replaced  by  gold  from  a  solution 
of  auric  chlorid  (AuCl3),  or  by  platinum  from  a  salt  of 
that  metal.  The  unchanged  silver  chlorid  is  removed  by 
sodium  thiosulfate  solution,  just  as  was  done  in  the  nega- 
tive. The  picture  is  then  washed  and  mounted. 

/ 

GOLD 

History  and  Occurrence.  Gold  was  considered  by  the  al- 
chemists to  be  the  most  perfect  metal,  and  many  spent  their 
lives  vainly  trying  to  transmute  the  baser  metals  into  it. 

Gold  is  usually  found  native.  The  most  famous  gold 
fields  have  been  those  of  California,  Australia,  South  Africa, 
and  Alaska.  It  usually  occurs  in  alluvial  deposits,  or  else 
in  quartz  veins. 

Extraction.  The  auriferous  gravel  or  clay  is  washed;  the 
lighter  materials  are  thus  allowed  to  be  carried  away,  while 
the  heavy  gold  remains  behind.  Gold-bearing  quartz  is 
crushed,  then  mixed  with  mercury,  and  amalgam  of  gold  is 


222        AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

formed.  The  mereury  is  distilled  off,  leaving  the  gold. 
Native  gold  usually  contains  silver.  After  both  metals  are 
dissolved  in  aqua  regia  (see  Exp.  61  c,  Part  I.)  the  silver  is 
precipitated  by  hydrochloric  acid.  The  auric  chlorid  is  then 
treated  with  ferrous  sulfate,  giving  metallic  gold,  which  is 
collected  and  melted  into  ingots.  Both  gold  and  silver  are 
frequently  reduced  in  blast  furnaces  in  connection  with  ores 
of  lead,  the  precious  metals  being  concentrated  in  the  lead 
ingot  formed.  When  gold  and  silver  are  mainly  associated 
with  copper,  they  are  obtained  in  the  copper  matte,  which  is 
produced  by  a  somewhat  complex  process  of  reduction  in 
the  reverberatory  furnace. 

Uses.  Coinage,  jewelry,  and  plate  consume  most  of  the 
gold  supply.  It  is  also  used  in  making  gold  leaf,  and  in 
making  ruby  glass. 

Auric  chlorid  is  obtained  by  dissolving  gold  in  aqua 
regia,  and  evaporating  carefully.  The  deliquescent  salt  that 
is  formed  is  used  extensively  in  photography. 

Other  Gold  Compounds  are  the  oxid  (Au2O3),  the  sulfid 
(Au2S3),  and  a  remarkably  beautiful  compound  called  the 
Purple  of  Cassius,  which  is  probably  a  double  stannate  of 
gold  and  tin. 

PLATINUM 

History.  Platinum  was  first  noticed  in  the  sixteenth 
century  by  Scaliger,  but  not  until  two  hundred  years  later 
did  it  become  well  known.  It  is  found  principally  in  the 
Ural  mountains,  California,  Mexico,  and  Australia. 

Extraction.  Platinum  is  extracted  from  its  ores  by  being 
treated  with  aqua  regia.  This  dissolves  the  platinum  to- 
gether with  certain  rare  metals.  After  the  solution  has  been 
evaporated  to  dryness,  the  rarer  chlorids  are  decomposed  by 
heat.  The  platinum  chlorid  is  then  dissolved  in  water,  pre- 


HISTORY,    OCCURRENCE    AND    APPLICATIONS         223 

cipitated  with  ammonium  chlorid,  and  heated.  A  spongy 
mass  of  metallic  platinum  results.  This  is  then  fused  to  a 
button  by  the  oxy-hydrogen  blowpipe. 

Uses.  Platinum  is  especially  valuable  for  crucibles, 
dishes,  and  other  utensils  of  the  chemical  laboratory.  It 
also  has  numerous  uses  in  electricity. 


QUALITATIVE  ANALYSIS 


QUALITATIVE   ANALYSIS 


THE  term  qualitative  analysis  signifies  the  art  of  finding 
out  what  elements  are  present  in  any  given  chemical  com- 
pound or  mixture.  Qualitative  analysis  is  not  concerned 
with  the  proportions  in  which  these  elements  may  be  present 
in  the  compound.  Such  determinations  belong  to  quantitative 
analysis.  Qualitative  analysis  has  for  its  basis  the  fact  that 
the  elements  fall  into  natural  groups  when  considered  with 
respect  to  the  compounds  they  form.  Thus  lead,  silver,  and 
mercurous  salts  are  naturally  grouped  together,  because, 
from  solutions  of  their  soluble  salts,  hydrochloric  acid  pre- 
cipitates them  as  chlorids.  Similarly  we  may  group  to- 
gether the  elements  that  are  precipitated  by  other  reagents. 
By  experiment  we  may  find  out  what  precipitate  is  obtained 
by  any  given  reagent  from  any  given  solution  of  a  salt  of  a 
single  metal.  Similarly,  we  may  find  solvents  that  will  sepa- 
rate precipitates  which  have  been  formed  by  a  reagent  acting 
upon  a  solution  containing  more  than  one  metal.  In  like 
manner  we  may  discover  the  precipitates  formed  by  reagents 
with  the  salts  of  any  acid.  Besides  precipitates,  we  may  use 
other  tests,  such  as  flame  colorations,  to  distinguish  the  ele- 
ments or  acid  radicals  from  one  another. 

The  following  table  gives,  with  their  solubilities,  the  im- 
portant precipitates  of  the  various  elements. 

227 


228      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


Special 


K,Cr04 


KCN 


fc  ^^ 


Na,HPO4 


(NH4)2S 


CD 


H2S 


Il 


H2S04 


53 

O  d 


HO 


NH4)2CO3 


Na.2CO3 


NH4OH 


2Q    • 

u'1 

o  " 


x'O 


NaOH 


QUALITATIVE    ANALYSIS 


229 


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S- 


61  a 


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23O      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


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QUALITATIVE    ANALYSIS 


231 


By  consulting  the  foregoing  tables,  it  is  possible  to  arrange 
the  compounds  of  the  metals  in  groups  according  to  their 
similarities,  thus :  — 


Lead 
Silver 
Mercurous 

Mercuric 

Copper 

Bismuth 

Cadmium 

Antimony 

Arsenic 

Tin 


Iron 

Aluminum 
Chromium 

Nickel 
Cobalt 
Manganese 
Zinc 


Barium 

Calcium 

Strontium 


GROUP  I. 
Precipitated  by  hydrochloric  acid. 


Precipitated  by  hydrogen  sulfid. 

GROUP  ill. 
Precipitated  by  ammonium  hydroxid. 


Precipitated  by  ammonium  sulfid. 

GROUP  v. 
Precipitated  by  ammonium  carbonate. 


GROUP    VI. 

Magnesium.     Precipitated  by  acid  sodium  phosphate. 


GROUP    VII. 

Sodium 

Potassium 

Lithium 

Ammonium.     Recognized  by  special  tests. 


Recognized  by  flame  colorations. 


232       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

How  to  Analyze  an  Inorganic  Substance 
EXAMINATION  FOR  BASES 

If  the  substance  is  not  in  solution,  it  may  be  dissolved  in 
water,  in  acids,  or,  after  fusing  with  sodium  carbonate,  it 
may  be  dissolved  in  acids. 

To  the  liquid  containing  the  mixture  to  be  analyzed,  add 
dilute  hydrochloric  acid.  Allow  the  precipitate,  if  one  is 
formed,  to  settle ;  then  add  a  drop  or  so  of  the  acid,  to  make 
sure  that  no  further  precipitate  can  be  formed.  Filter,  wash 
with  cold  water,  and  save  the  filtrate.  The  precipitate  con- 
tains Group  I. 

GROUP    I. 

Break  a  hole  in  the  filter,  and  wash  the  precipitate  into  a 
test  tube  by  means  of  a  fine  stream  of  water  from  a  wash 
bottle.  Boil  and  filter  again.  The  filtrate  from  this  con- 
tains the  soluble  lead  chlorid,  if  any  is  present.  This  may 
be  tested  for  lead  by  potassium  chromate,  which  gives  a  yellow 
precipitate  soluble  in  sodium  hydroxid. 

To  the  precipitate  (if  there  is  any)  remaining  on  the  filter, 
add  ammonium  hydroxid,  and  collect  the  resulting  filtrate  in 
a  beaker.  If  the  precipitate  turns  black,  a  mercurous  com- 
pound is  present.  To  the  filtrate  add  nitric  acid.  If  any 
silver  chlorid  was  dissolved,  it  will  be  reprecipitated  by  the 
nitric  acid,  thus  indicating  the  presence  of  silver. 

The  filtrate  from  Group  I.  may  contain  the  compounds  of 
all  elements  belonging  to  the  succeeding  groups.  Dilute 
this  with  water,  and  heat  to  boiling;  then,  under  a  hood, 
pass  a  stream  of  hydrogen  sulfid  through  the  solution  until  it 
is  saturated.  The  solution  is  saturated,  if,  on  removing  the 
beaker  and  blowing  away  the  gas  from  above  the  liquid,  it 


QUALITATIVE    ANALYSIS  233 

still  smells  of  hydrogen  sulfid  after  a  lapse  of  one  or  two 
minutes.  Then  filter,  and  save  the  filtrate.  Wash  the  pre- 
cipitate until  the  wash  water  is  no  longer  acid.  The  pre- 
cipitate contains  Group  IT. 

GROUP    II. 

Remove  the  precipitate  to  a  porcelain  dish,  add  enough 
concentrated  ammonia  to  cover  the  mixed  sulfids,  and  warm, 
stirring  the  mixture  continually.  Then  add  a  little  yellow 
ammonium  sulfid  together  with  the  ordinary  kind,  until  the 
mixture  smells  distinctly  of  the  sulfid.  After  warming  a  few 
minutes,  allow  the  mixture  to  stand  about  fifteen  minutes. 
Then  filter,  and  wash  the  precipitate  with  hot  water.  The 
precipitate  may  contain  any  or  all  of  A.  of  Group  II.,  i.e., 
mercuric,  copper,  bismuth,  or  cadmium  sulfids.  The  filtrate 
may  contain  any  or  all  of  B.  of  Group  II.,  i.e.,  antimony, 
arsenic,  or  tin  sulfids. 

A.    OF    GROUP    II. 

Remove  the  precipitate  to  a  porcelain  dish,  and  add 
enough  dilute  nitric  acid  to  cover  the  mixed  sulfids  about 
twice  over.  Boil,  and  stir  the  mixture  continually  until  all 
solvent  action  has  ceased.  Filter  and  wash.  The  precipi- 
tate contains  mercuric  sulfid  and  any  lead  that  may  have 
been  dissolved  by  the  water  from  Group  I.  The  presence  of 
the  mercuric  sulfid  is  confirmed  by  adding  a  few  drops  of 
aqua  regia  to  the  precipitate,  diluting,  and  then  adding  a 
little  stannous  chlorid.  If  a  white  precipitate,  which  blackens 
on  the  addition  of  ammonia,  is  formed,  the  indication  is  that 
mercury  is  present. 

The  filtrate  may  contain  copper,  bismuth,  and  cadmium 
nitrates.  Add  ammonium  hydroxid  to  excess.  A  white  floc- 
culent  precipitate  indicates  bismuth,  and  a  blue  color  in  the 


234      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

solution  indicates  copper.  Filter,  and  add  potassium  cyanid 
until  the  blue  color  disappears ;  then  pass  a  current  of  hydro- 
gen sulfid  through  the  solution.  A  yellow  precipitate  indi- 
cates cadmium. 

B.    OF    GROUP    II. 

To  the  filtrate  obtained  from  the  mixed  sulfids  in  A.,  add 
dilute  hydrochloric  acid.  If  the  precipitate  is  milky,  probably 
nothing  but  sulfur,  formed  by  the  decomposition  of  the 
alkaline  sulfid,  is  present.  If  the  precipitate  is  lemon  yellow, 
or  orange,  at  least  one  sulfid  of  the  metals  of  B.  is  present. 
Filter,  remove  the  precipitate  to  a  porcelain  dish,  and  boil 
it  with  concentrated  hydrochloric  acid.  Hydrochloric  acid 
decomposes  the  sulfids  of  antimony  and  ##,  leaving  the  arsenic 
sulfid.  Filter.  The  precipitate  remaining  may  be  confirmed 
as  arsenic  by  removing  it  to  a  porcelain  dish,  and  adding 
concentrated  hydrochloric  acid and  a  few  crystals  of  potassium 
chlorate.  Evaporate  off  the  excess  of  acid ;  then  add  ammo- 
nium chlorid,  an  excess  of  ammonium  hydroxid,  and  mag- 
nesium sulfate.  If  arsenic  is  present,  needle-shaped  crystals 
of  ammonium  magnesium  arsenate  will  be  deposited. 

Dilute  the  filtrate  with  water,  and  add  a  strip  of  zinc  and 
a  strip  of  platinum.  A  black  coating  upon  the  platinum 
indicates  antimony.  If  tin  is  present,  a  gray  powder,  partly 
on  the  bottom  of  the  beaker  and  partly  on  the  zinc,  is  formed. 
This  powder  may  be  removed  and  dissolved  in  concentrated 
hydrochloric  acid.  Mercuric  chlorid,  added  to  the  solution, 
gives  a  white  precipitate  of  mercurous  chlorid.  This  precipi- 
tate soon  becomes  grayish,  owing  to  its  reduction  to  metallic 
mercury. 

The  filtrate  from  Group  II.  may  contain  compounds  of 
elements  belonging  to  the  succeeding  groups.  Boil  the  fil- 
trate until  all  traces  of  hydrogen  sulfid  have  been  removed. 


QUALITATIVE    ANALYSIS  235 

Then  add  a  few  drops  of  nitric  add,  to  change  any  ferrous 
compounds  present  to  ferric.  To  the  hot  solution,  add  a 
moderate  quantity  of  ammonium  chlorid.  While  stirring  the 
solution,  add  ammonium  hydroxid*  in  small  quantities  until 
a  distinct  odor  of  ammonia  is  perceptible.  (The  ammo- 
nium chlorid  is  added,  to  keep  in  solution  the  compounds  of 
metals  belonging  to  succeeding  groups.)  Filter  and  wash. 
Save  the  filtrate.  The  precipitate  contains  Group  III. 

GROUP    III. 

If  the  precipitate  is  white,  only  aluminum  is  present ;  if  it 
is  grayish-green  or  blue,  chromium  is  present ;  if  it  is  brown, 
iron  is  present.  Dissolve  the  precipitate,  on  the  filter,  by 
adding  dilute  hydrochloric  acid  drop  by  drop.  Dilute  the 
resulting  solution  with  water,  add  sodium  hydroxid  to  excess, 
and  boil.  The  chromium  and  iron  are  thus  precipitated, 
while  the  aluminum  hydroxid  formed  is  dissolved  by  the 
alkali.  When  the  precipitate  has  settled,  filter.  The  pre- 
cipitate may  be  examined  for  chromium  by  fusing  a  small 
quantity  with  a  mixture  of  potassium  nitrate  and  sodium  car- 
bonate upon  a  piece  of  platinum  foil,  dissolving  the  mass 
thus  obtained  in  water,  adding  a  little  acetic  acid  to  liberate 
the  carbon  dioxid,  and  then  adding  a  little  lead  acetate.  A 
yellow  precipitate  indicates  chromium.  Iron  in  the  precipi- 
tate from  sodium  hydroxid  is  shown  by  dissolving  a  little  of  it 
in  hydrochloric  acid  and  testing  it  with  potassium  ferrocyanid. 

The  filtrate  containing  the  excess  of  sodium  hydroxid  is 
then  tested  for  aluminum  by  acidifying  with  hydrochloric  acid, 
and  adding  ammonium  hydroxid.  A  white  flocculent  pre- 
cipitate indicates  aluminum. 

*  It  is  assumed  that  there  are  no  phosphates  or  oxalates  present  in 
the  solution.  Their  presence  would  make  necessary  a  different  treat- 
ment of  the  solution, 


236      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

The  filtrate  obtained  from  Group  III.  may  contain  com- 
pounds of  elements  belonging  to  succeeding  groups.  Trans- 
fer the  solution  to  a  flask,  and  heat  it  to  boiling.  Add 
ammonium  sulfid  to  the  hot  solution,  drop  by  drop,  as  long 
as  any  precipitate  is  formed,  shaking  the  flask  continually. 
Allow  the  precipitate  to  settle,  filter,  and  wash.  Save  the 
filtrate.  The  precipitate  may  contain  any  or  all  of  the 
sulfids  of  nickel,  cobalt,  manganese,  and  zinc,  i.e.,  Group  IV. 

GROUP    IV. 

Place  the  precipitate,  paper  and  all,  in  a  porcelain  dish, 
and  add  cold  dilute  hydrochloric  acid.  This  decomposes  the 
sulfids  of  manganese  and  zinc,  and  leaves  the  sulfids  of  nickel 
and  cobalt  unchanged.  Filter  and  wash.  Treat  the  precipi- 
tate with  a  little  aqua  regia,  and  evaporate  it  to  dryness. 
Dissolve  the  resulting  chlorids  in  water,  and  add  a  medium 
sized  piece  of  potassium  nitrite.  Yellow  crystals  si  potassium 
cobalt  nitrite  will  form,  if  cobalt  is  present.  Filter  these  off, 
and  test  the  filtrate  for  nickel  by  adding  sodium  hydroxid. 

The  filtrate,  obtained  by  treating  the  original  mixed  sulfids 
with  dilute  hydrochloric  acid,  is  now  boiled  until  all  hydrogen 
sulfid  is  expelled.  Sodium  hydroxid  is  then  added,  and,  if 
manganese  is  present,  a  white  gelatinous  precipitate  of  man- 
ganese hydroxid  is  formed.  The  precipitate  rapidly  turns 
brown.  The  filtrate  from  this  contains  the  zinc  hydroxid,  if 
any  is  present,  dissolved  in  excess  of  sodium  hydroxid. 
Pass  hydrogen  sulfid  through  this.  A  white  precipitate, 
soluble  in  hydrochloric  acid,  indicates  zinc. 

The  filtrate  obtained  from  the  precipitate  of  Group  IV. 
may  contain  any  or  all  of  the  compounds  of  metals  belong- 
ing to  the  succeeding  groups.  Boil  the  filtrate  until  all 
ammonium  sulfid  is  decomposed,  and  filter.  Then  add  a 
little  ammonium  chlorid,  after  which  add  ammonium  carbonate 


QUALITATIVE    ANALYSIS  237 

as  long  as  any  precipitate  is  formed ;  then  warm  gently. 
The  precipitate  may  contain  any  or  all  of  the  carbonates 
of  Group  V.  Filter  and  wash.  Save  the  filtrate. 

GROUP  v. 

Wash  the  precipitate  into  the  apex  of  the  filter,  and  then 
add  acetic  acid,  catching  in  a  test  tube  the  solution  that  runs 
through.  To  this  add  potassium  chromate  ;  barium  chromate 
is  formed,  if  barium  is  present.  Filter,  and  to  the  filtrate 
add  dilute  ammonium  hydroxid.  Then  add  ammonium  car- 
bonate to  reprecipitate  the  calcium  and  strontium,  and  filter 
again.  Wash  until  the  precipitate  is  white,  and  then  redis- 
solve  it  in  acetic  acid.  Add  to  this  a  dilute  solution  of 
ammonium  sulfate  (i  part  sulfate  to  12  parts  of  water),  and 
allow  the  precipitate,  if  one  is  formed,  to  settle.  A  white 
precipitate  indicates  strontium.  Then  filter,  and  add  a  little 
acetic  acid  and  ammonium  oxalate.  A  white  precipitate, 
soluble  in  hydrochloric  acid  and  reprecipitated  by  ammonium 
hydroxid,  indicates  calcium. 

The  filtrate  obtained  from  the  precipitation  of  Group  V. 
may  contain  the  compounds  of  metals  belonging  to  Groups 
VI.  and  VII.  The  reagent  for  Group  VI.  is  acid  sodium 
phosphate.  Since  this  reagent  contains  sodium,  it  would  evi- 
dently be  useless  to  test  for  sodium  in  the  original  solution, 
after  once  having  added  the  reagent.  It  is  hence  necessary 
to  divide  the  filtrate  from  Group  V.  into  two  parts,  and  ex- 
amine one  part  for  Group  VI.,  and  the  other  for  Group 
VII. 

GROUP    VI. 

Place  one  portion  of  the  filtrate  from  Group  V.  in  a  large 
test  tube,  and  add  ammonium  hydroxid  and  acid  sodium 
phosphate.  Shake  the  solution  well,  scratch  the  inside  of 


238      AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

the  tube  with  a  glass  rod,  and  allow  the  mixture  to  stand. 
A  white  crystalline  precipitate  indicates  magnesium. 

GROUP    VII. 

The  remaining  part  of  the  filtrate  from  Group  V.  is  evapo- 
rated to  dryness,  and  is  then  heated  so  that  all  ammonium 
compounds  are  decomposed.  It  may  then  be  tested  for 
potassium  and  sodium  by  the  color  that  the  residue  gives  to 
a  non-luminous  flame.  If  sodium  is  present,  the  character- 
istic sodium  flame  will  be  seen.  This  flame,  however,  will 
obscure  any  coloration  due  to  potassium ;  therefore  it  is 
necessary  to  view  the  flame  through  a  piece  of  cobalt  glass. 
The  cobalt  glass  shuts  off  the  yellow  sodium  rays,  and 
allows  the  violet  potassium  rays  to  pass  through.  Lithium 
gives  a  beautiful  red  color  to  the  flame.  In  order  to  test 
for  the  presence  of  ammonium  compounds,  it  is  necessary  to 
use  the  original  mixture.  To  a  small  test  of  this,  add  sodium 
hydroxid.  If,  on  boiling,  the  steam  turns  red  litmus  paper 
blue,  ammonium  is  present. 

EXAMINATION    FOR    ACIDS 

In  examining  a  solution  for  acid  radicals,  we  do  not  have 
the  advantage  of  such  a  well-defined  and  systematic  scheme 
as  we  have  in  examining  it  for  bases.  We  are,  in  this  case, 
obliged  to  depend  upon  a  knowledge  of  the  relations  that 
various  acids  bear  to  various  bases,  and  also  upon  numerous 
special  tests.  For  instance,  if  we  find  a  lead,  silver,  or 
mercurous  salt  in  a  solution,  we  know  that  no  halogen 
can  be  present ;  for  the  halogen  would  then  precipitate  the 
metal.  Likewise,  no  sulfate  can  be  found  in  a  solution 
containing  a  salt  of  barium. 

The  special  tests  for  the  various  common  acids  may  be 
found  by  referring  to  the  following  table, 


QUALITATIVE    ANALYSIS 


239 


* 

1 

C                                            0* 

£                         < 

6C                                   .C 

<                    a. 

2 
^ 
be 
X 

i 

HC1 

w. 

sol. 
NH4OH 

W. 

W. 
turns  Blk. 
with  NH4OH 

HBr 

Yellowish                W. 
sol.                     so). 
KCN             dil.  HNOS 

Y.  W. 

HI 

Y. 

sol. 
KCN 

V. 
sol. 
dil.  HNO3 

Gr.  Y. 

sol. 
KI 

HgCXO.Jj 

gives 
R. 
sol.  KI 

H,S03 

W. 

W. 

sol. 
dil.  HNO3 

w. 

sol. 

dil.  HC1 

H,S04 

W. 

sol. 
tartaric 
acid 
and 
NH4OH 

W. 

insol  . 
all  dil. 
acids 

4 

H,P04 

V. 
sol. 
NH.OH, 
HNOa 

W 
sol. 
HN03 

W. 

sol. 
HC1.HXO, 

HCN 

W 
sol. 
NH4OH 

Add  small  amount  FeCl3  and  FeSO4;    make  alkaline 
with  NH4OH;  warm  and  acidify  with  HC1.    Turns  blue. 

HF 

Gives  HF  with  H,SO4. 

H,S 

Various  precipitates  with  metals. 

HN03 

Two  or  three  cc.  cone.  H2SO4,  poured  carefully  to  bottom  of  sol.  in 
test  tube  and  FeSO4  sol.  added,  gives  brown  ring. 

H3B03 

Colors  alcohol  flame   green. 

H,COa 

Gives   CO2   with   acids. 

H4Si  O4 

Give  insoluble  skeleton  in  microcosmic  bead. 

240       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 


APPENDIX    I. 
Metric  System 

LINEAR    MEASURE 


10  millimeters 
10  centimeters 
10  decimeters 
10  meters 
10  dekameters 
10  hektometers 


i  centimeter 
i  decimeter 
i  meter 
i  dekameter 
i  hektometer 
i  kilometer 


CAPACITY 


100  centiliters 
100  liters 


liter 
hektoliter 


i  liter      =      i  cubic  decimeter 

=      1000  cubic  centimeters 


WEIGHT 


10  milligrams 

=      i  centigram 

10  centigrams 

=      i  decigram 

10  decigrams 

==      i  gram 

10  grams 

=      i  dekagram 

10  dekagrams 

=      i  hektogram 

10  hektograms 

—      i  kilogram 

oo  milligrams 

—      i  gram 

oo  grams 

=      i  kilogram 

oo  kilograms 

=      i  tonneau 

i  cubic  centimeter  of  water  at  4°  C.  weighs  i  gram, 
i  cubic  decimeter  of  water  at  4°  C.  weighs  i  kilogram. 


APPENDIX    II.  241 

APPENDIX    II. 
Manipulations 

1 .  To  Cut  Glass  Tubes.     Mark  the  glass  with  a  file  at  the 
required  point,  then,  seizing  the  tube  with  the  hands  so  that 
the  thumbs    are  opposite  the  mark,  break  by  pulling  with 
the  hands  and  at  the  same  time  pressing  outward  with  the 
thumbs. 

To  cut  large  glass  tubes,  make  a  mark  with  a  file  as  be- 
fore. Take  two  pieces  of  filter  paper  long  enough  to  go 
around  the  tube  several  times,  and  fold  them  twice  so  as  to 
form  a  strip  of  four  thicknesses,  3  cm.  wide.  After  wetting 
the  papers  thoroughly,  wrap  them  around  the  tube  about 
4  mm.  from  the  mark,  one  on  each  side  of  it.  Be  careful  to 
have  the  edges  of  the  wet  filter  paper  straight  and  touching 
the  glass  at  all  points.  Then  rotate  the  wrapped  portion  of 
the  tube  in  the  hot  flame  of  a  Bunsen  burner,  when,  after  a 
few  moments,  the  tube  will  break  off  evenly. 

2.  To  Bend  Glass  Tubes.     A  fish-tail  gas  flame  should  be 
used.     Hold  the  tube  lengthwise  in  the  flame  just  above  the 
dark   part,  at  the  same    time   rotating  it  with  the   fingers. 
When    the    glass    has    softened   enough,  remme  it  from  the 

flame  and  quickly  bend  it  to  the  desired  angle.     Allow  the  tube 
to  cool  with  the  soot  upon  it. 

3.  To  Insert  a    Glass  Tube  into  a  Rubber  Stopper.     Wet 
both  glass  and  hole.     The  tube  can  then  be  easily  pushed 
in  the  required  distance.     Be  very  careful  in   pushing  tubes 
into  rubber  stoppers,  since  a  glass  tube  broken  while  held 
in    the    hand    in    this    position    sometimes    makes    an    ugly 
wound. 

4.  To  Bore  a  Hole  through  Glass.     Procure  a  rat-tail  file 
about  4  mm.   in  diameter.     Break  the  file   and  grind   the 


242       AN    ELEMENTARY    EXPERIMENTAL    CHEMISTRY 

broken  end  down  on  two  sides  to  a  convex  edge.  Moisten 
the  file  and  the  glass  with  the  following  solution :  — 

i  oz.  Camphor  Gum, 
ii  oz.  Oil  of  Turpentine. 

Directly  under  the  point  where  the  hole  is  desired,  place  a 
small  piece  of  hard  wood  about  i  sq.  cm.  on  the  base  with 
an  upper  surface  of  about  4  sq.  mm.  Apply  the  end  of  the 
file,  using  it  as  you  would  an  awl.  When  the  glass  is  pierced, 
grind  the  hole  out  to  the  required  size  by  means  of  a  smaller 
rat-tail  file.  Moisten  the  file  occasionally  with  the  solution. 

5.  To  Dry  a  Flask.     First  rinse   the  flask  with  a   little 
alcohol.     Take  a  tube  several  inches  longer  than  the  flask, 
and  insert  it  to  the  bottom.     Then  force  air  through  the  tube 
with  a  bellows,  at  the   same   time   rotating  the  tube   in  the 
flame  of  a  Bunsen  burner.     The  alcohol  will  be  vaporized, 
and  the  flask  will  be  dry  in  a  few  moments. 

6.  To  Make  a  Wash  Bottle.     Fit  a  two-holed  rubber  stopper 
with  two  glass  tubes.     Let  one  reach  to  the  bottom  of  a  liter 
flask,  and  have  its  outer  end  bent  at  an  angle  of  45°.     The 
outer  arm  of   this   tube  should  be  about  3   cm.   long,  and 
should  be  connected  by  means  of  rubber  tubing  with  a  short 
glass  tube  of  the  same  bore  drawn  out  to  a  diameter  of  about 
i   mm.     The  second  tube  should  extend  only  through  the 
stopper,  and  its  outer  portion  should  be  bent  at  an  angle  of 
135°.     Round  off  the  ends  of  both  tubes  in  the  Bunsen  flame. 
To  make  the  flask  easy  to  handle  when  containing  hot  water, 
bind  a  strip  of  cork  around  its  neck. 

7.  Filters.     Filter   papers  are   usually  folded,  first  along 
one  diameter  and  then  along  a  line  perpendicular  to  this. 
When  the  filter  is  opened,  one  side  has  three  thicknesses  of 
paper,  and  the  other  side  has  one. 

8.  To  Use  a  Blowpipe.     The  use  of  the  blowpipe  requires 
a  constant  passage  of  a  stream  of  air  through  it.     To  pro- 


APPENDIX    II.  243 

duce  this,  fill  the  mouth  with  air,  at  the  same  time  breathing 
through  the  nose.  The  muscular  action  of  the  cheeks  will 
force  a  constant  current  of  air  through  the  blowpipe.  Do 
not  attempt  to  breathe  through  the  mouth,  but  replenish  the 
air  in  the  mouth  from  the  lungs.  A  little  practice  will  make 
you  so  expert  that  you  will  be  able  to  blow  a  constant  flame 
for  many  minutes. 


INDEX 


Absolute  zero,  103. 

Acetylene,  189. 

Acid,  definition  of,  12,  29. 

arsenious,  74. 

arsenic,  74. 

boric,  185. 

carbonic,  12. 

hydriodic,  44,  169. 

hydrobromic,  42,  167. 

hydrochloric,  35,  36,  37,  38,  164. 

hydrocyanic,  192. 

hydrofluoric,  45,  169. 

hydrosulfuric,  47,  171. 

muriatic  (see  hydrochloric  acid). 

nitric,  59,  62,  179. 

nitrous.  178. 

oxalic,  50. 

phosphoric,  u,  183. 

phosphorous,  n. 

prussic,  192. 

sulfuric,  16,  173. 

sulfurous,  14. 
Acids,  examination  of  substances 

for,  238, 
Air,  a  mechanical  mixture,  57. 

density  of,  107. 

relation  to  combustion  and  de- 
cay, 58. 

weight  of  a  liter  of,  104. 
Alcohol,  155. 
Alkali,  1 8. 


Allotropic  forms,  10. 
Aluminum,  79,  206. 

action  of  acids  on,  80. 

action  upon  alkalis,  80. 

alloy,  217. 
Alums,  207. 

Amalgamation  process,  219. 
Amalgams,  218. 
Ammonia,  66,  177. 

from  ammonium  chlorid,  68. 

from  organic  matter,  67. 
Ammonium,  69. 

amalgam,  69. 

chlorid,  67. 

hydroxid,  69. 
Analysis,  7. 
Animal  charcoal,  187. 
Antichlor,  166. 
Antimony,  74,  184. 

chlorid,  75. 

hydrid,  75. 

mirror,  74,  75. 

oxid,  75. 
Aqua  regia,  84. 
Argon,  181. 
Arsenic,  72. 

mirror,  72,  74. 

oxid,  72. 

reduction  of,  73, 
Arsenic  acid,  74.      , 
Arsenious  acid,  74. 
Arsin,  73. 


245 


246 


INDEX 


Atmospheric  pressure,  97. 
Atomic  theory,  89,  96. 
Atomic  weights,  108,  124,  i: 
Atoms,  97. 

Avogadro's  hypothesis,  120. 
Azote,  177. 


Bacilli,  58. 

Baking  powder,  201. 

Barium,  81,  204. 

chlorid,  81. 

flame  test,  82. 

hydro xid  and  oxid,  81. 

nitrate,  Si. 

sulfate,  82. 
Barometer,  97,  98. 
Bases,  18,  29. 

examination  for,  232. 
Basicity,  135. 
Benzine,  188. 
Binary,  137. 
Bismuth,  75,  214. 

basic  nitrate,  76. 

nitrate,  75. 

oxid,  75. 

sulfid,  76. 

Bituminous -coal,  187. 
Blast  furnace,  209. 
Bleaching  powder,  166. 
Blowpipe,  242. 
Borax,  186. 
Boric  acid,  185. 
Boron,  185. 
Boyle's  law,  99,  100. 
Brass,  205. 
Bread,  20 T. 
Bricks,  195. 
Britannia  metal,  185. 


Bromin,  41,  167. 
action    upon    hydrogen    sulfid, 

49- 

replacement  of,  by  chlorin,  43. 
Bronze,  214. 


Cadmium,  76,  206. 

action  of  acids  upon,  76. 

oxid,  76. 

sulfid,  77. 
Calcium,  30,  202. 

carbonate,  31,32,  202. 

chlorid,  39,  204. 

fluorid,  45. 

hydroxid,  31,  203. 

oxid,  31,  203. 

sulfate,  203. 
Calorie,  113. 
Carbon,  12,  186. 

bisulfid,  192. 

dioxid,  39,  192. 
density  of,  107. 

relation  to  combustion  and  de- 
cay, 58. 

relation  to  growing  plants,  58. 
weight  of  liter  of,  105. 

monoxid,  49,  52,  191. 
Carbonic  acid,  12. 
Cements,  203. 
Cells,  59. 

Changes,  chemical  and  physical,  i. 
Charcoal,  12,  187. 
Charles,  law  of,  101. 
Chemical  compound,  2. 

equations,  138-140. 

equivalence,  109. 

symbols,  127,  128. 
Chlorin,  34,  40,  163. 


INDEX 


247 


Chlorin,  acids   and   oxy-acids   of, 
165. 

and  hydrogen,  35. 
Chromium,  208. 
Clay,  195. 
Coal,  1 86. 
Coal  tar,  189. 
Cobalt,  213. 
Coke,  187. 
Combustion,  n. 
Condenser,  155. 
Copper,  5,  216. 

oxid,  5. 

sulfid,  48. 

sulfate,  217. 
Cupellation  process,  219. 

D 

Daguerreotype,  220. 

Dalton's  hypothesis,  96. 

Definite  proportions  by  weight,  92. 

by  volume,  119,  120. 
Density,  107. 
Developing  pictures,  220. 
Diamond,  107. 

Disinfectant,  163, 164,166, 172, 173. 
Dissociation,  151,  152,  153. 
Divalent,  132. 
Dolomite,  202. 
Dry  distillation,  154. 
Dulong  and  Petit,  law  of,  117,  n  8. 
Dyad,  132. 

£ 

Efflorescence,  28. 

Electrical  equivalents,  112,  113. 

Electrolysis,  8. 

Elements,  3. 

electro,  relations  of,  137. 
Emerald,  208. 


Endothermic  actions,  145. 
Energy,  potential,   kinetic^  chemi- 
cal, 145. 

Epsom  salt,  204. 
Equations,  158-140. 
Etching  glass,  46. 
Ethylene,  53,  189. 
Exothermic  actions,  145. 

F 

Factors,  90. 
Fermentation,  155. 
Fertilizers,  204. 

Ferrous  and  ferric  chlorids,  So. 
Ferrous  sulfate,  212. 
Fire  damp,  188. 
Fixing  pictures,  221. 
Flame,     effect     of    excluding    air 
from,  7. 

nature  of,  54. 
Fluorin,  45,  169. 

Formula  for  changes  of  tempera- 
ture and  pressure,  103. 
\    Furnace,  blast,  209. 

reverberatory,  199. 

G 

Galvanized  iron,  205. 
Gas  carbon,  187. 
Gas,  illuminating,  189. 

natural,  188. 

water,  191. 
Gasolene,  188. 
Gay  Lussac,  law  of,  120. 

tower,  17.4. 
German  silver,  213. 
Germs,  58. 
Glass,  193. 

Glauber's    salt    (sodium    sulfate) 
199. 


248 


INDEX 


Glover  tower,  174. 
Glycerin,  180. 
Gold>.84,  221. 

and  aqua  regia,  84. 

chlorid,  84. 

Graphic  formulae,  132. 
Graphite,  12,  186. 
Gun  cotton,  180. 
Gunpowder,  197. 
Gypsum,  203. 

H 

Haemoglobin,  58. 
Halogens,  163. 
Heat  unit,  1 13. 

specific,  113. 

of  chemical  action,  147. 

of  neutralization,  149. 

of  solution  and"  of  hydration,  1 50. 
Helium,  181. 
Hydraulic  main,  190. 
Hydriodic  acid,  44,  169. 
Hydrobromic  acid,  42,  167. 
Hydrochloric   acid,  35,  36,  37,  38, 

164. 

Hydrocyanic  acid,  192. 
Hydrofluoric  acid,  45,  169. 
Hydrogen,  160. 

antimonid,  75. 

arsenid,  73. 

burning  of,  in  chlorin,  35. 

from  iron  and  sulfuric  acid,  25. 

from  potassium  and  water,  23. 

from  sodium  and  water,  22. 

from  zinc  and  sulphuric  acid,  23. 

sulfid,  47,  48,  171. 

action  of  bromin  upon,  49. 

union  with  chlorin  by  light,  35. 

weight  of  liter  of,  106. 


Hydrosulfuric  acid  (see  hydrogen 
sulfid). 

I 

Ice,  3. 

Illuminating  gas,  189. 
Ink,  212. 

Indestructibility  of  matter,  90. 
lodin,  43,  1 68. 

and  phosphorus,  43. 

extraction  from  seaweed,  168. 

replacement  by  bromin,  45. 

replacement  by  chlorin,  45. 
Ions,  153. 
Iron,  20.  • 

chlorids,  80. 

galvanized,  205. 

oxid,  20,  21. 

pig,  210,  211. 

specific  heat  of,  116. 

sulfate,  212. 

sulfid,  48. 

wrought,  210. 
Isomorphic  equivalents,  119. 


Kerosene,  188. 

Kindling  temperature,  13. 


Lamp  black,  187. 
Laughing  gas,  71."* 
Law,  definition  of,  90. 

of  Boyle,  99,  100. 

of  Charles,  101. 

of  definite  proportion  by  wreight, 
92. 

of  definite  proportions   by   vol- 
ume, 1 20. 

of  Dulong  and  Petit,  117,  118. 


INDEX 


249 


Law  of  Gay  Lussac,  120. 

of  indestructibility  of  matter,  90. 
Law  of  multiple  proportions,  93. 
Lead,  77,  215. 

acetate,  216. 

action  of  acids  upon,  77. 

chlorid,  78. 

oxids,  77. 

replacement  of,  by  zinc,  78. 

specific  heat  of,  114. 

sulfate,  78. 

sulfid,  78. 

white,  216. 

Le  Blanc  process,  200. 
Lime,  31,  203. 
Limewater,  31. 

Limewater  and  carbonic  acid,  31. 
Litharge,  216. 


Magnesium,  20,  204. 

oxid,  20. 

oxid  and  water,  20. 

oxid  and  sulfuric  acid,  28. 
Manganese,  212. 

test  for,  40. 

chlorid,  41. 

dioxid,  7. 

sulfate,  213. 
Marble,  analysis  of,  32. 
Marsh  gas,  52,  53,  188. 
Matches,  183. 
Matter,  three  conditions  of,  3. 

indestructibility  of,  90. 
Mechanical  mixture,  4. 
Mendele Jeff's  table,  126. 
Mercurous  chlorid,  77. 
Mercury,  6,  77,  217. 

and  oxygen,  6,  7. 


Microcrith,  118. 
Micro-organisms,  58. 
Molecular  weights,  120. 

formulae,  determination  of,  129. 
Molecules,  97. 
Monad,  132. 
Mordant,  207,  208. 
Mortar,  203. 

Muriatic    acid    (see    hydrochloric 
acid). 


Naming  of  compounds,  137. 

Naphtha,  188. 

Nascent  state  of  elements,  65. 

Natural  gas,  188. 

Nature  of  flame,  54. 

Negative,  electro,  137 

photographs,  220. 
Neutralization,  28. 
Nickel,  Si,  213. 

action  of  acids  upon,  81. 

electrical  equivalent  of,  112. 
Niter  plantations,  197. 
Nitric  acid,  59,  62,  179. 

instability  of,  62. 

neutralization  of,  62. 
Nitric  oxid,  63. 

a  reducing  agent,  71. 

peroxid,    an     oxidizing     agent, 

71- 

Nitrous  acid,  178. 
Nitrous  oxid,  70. 

analysis  of,  72. 
Nitrogen,  57,  177. 

acids  of,  178. 

oxids  of,  178. 
Nitroglycerin,  180. 
Normal  salt,  29. 


250 


INDEX 


Olefiant  gas,  53,  54. 
Organic  chemistry,  154. 
Organic  matter,  burning  of,  154. 
Oxalic  acid,  ana-ysis  of,  50. 
Oxidation,  27. 
Oxygen,  7,  159. 

atomic  weight  of,  122. 

burning  of,  in  hydrogen,  26. 

density  of,  122. 

molecular  weight  of,  123. 

number  of   atoms  to  the  mole- 
cule, 121. 

preparation     from      manganese 
dioxid,  7. 

from  potassium  chlorate,  8. 
by  electrolysis,  8. 
Oxy-hydrogen  blowpipe,  161. 
Ozone,  10. 

number  of   atoms   in  molecule, 
122. 

P 

Paraffin,  188. 
Paris  green,  217. 
Peat,  187. 

Percentage    composition,    calcula- 
tion of,  144. 

Periodicity  of  the  elements,  125. 
Petroleum,  188. 
Pewter,  216. 
Phosphin,  183. 
Phosphoric  acid,  II,  183. 
Phosphorous  acid,  n. 
Phosphorus,  10,  181. 

and  iodin,  43,  44. 

and  oxygen,  10,  11. 
Photography,  220. 
Physical  change,  3. 
Physiological  chemistry,  59. 


Pig  iron,  210,  211. 
Plants,  chemistry  of,  58. 
Plaster  of  Paris,  204. 
Platinum,  85,  222. 

action  of  acids  on,  85. 

action  of  hot  metals  on,  85. 
Porcelain,  194. 
Positive,  electro,  137. 
Positive  photographs,  220. 
Potassium,  18,  195. 

bromid,  43. 

and  sulfuric  acid,  43, 

carbonate,  196. 

chlorate,  8,  198. 

molecular  weight  of,  123. 

chlorid,  39. 

molecular  weight  of,  124. 

hydro  xid,  19,  196. 

ferrocyanid,  212. 

iodid,  44. 

and  sulfuric  acid,  44. 

and  phosphoric  acid,  44. 

oxid,  1 8. 
Pottery,  195. 
Precipitates,    tables    of,    228,  229, 

230,  239. 
Products,  90. 
Prout's  hypothesis,  127. 
Prussian  blue,  212. 
Prussic  acid,  192. 
Puddling,  210. 
Putreficadon,  58. 


Qualitative  analysis,  227. 
Quantitative  analysis,  94,  227. 
Quartz,  193. 
Quick  lime,  31,  203.     ^ 
Quicksilver  (see  mercury). 


INDEX 


251 


Radical,  69. 

Reactions,  writing  of,  138. 
Reduction,  27. 
Reverberatory  furnace,  199. 

S 

Safety  lamp,  55. 
Salt,  definition  of,  29. 

acid,  basic,  and  normal,  29. 

cake,  199. 
Saponification.  156. 
Seaweed,  168. 
Selenium,  176. 
Shot,  216. 
Silicon,  193. 

dioxid,  193. 
fluorid,  46. 
Silver,  83,  219. 

action  of  acids  upon,  83. 

bromid,  83. 

chlorid,  83. 

reduction  of,  by  hydrogen,  84. 
analysis  of,  94. 

iodid,  83. 

oxid,  83. 

replacement  of,  by  copper,  84. 
Soap,  156,  197. 
Sodium,  17,  198. 

amalgam,  18. 

carbonate,  200. 

chlorid,  36,  198. 
analysis  of,  94. 

hydro xid,  17,  201. 

and  sulfuric  acid,  28. 

oxid,  17. 

sulfate,  28,  29,  199. 
Solder,  205,  214,  216. 
Solvay  process,  200. 


Specific  heat,  113. 

Spiegeleisen,  211,  213. 

Steel,  211. 

Stibin,  75. 

Stoic  biometry,  141. 

Strontium,  82,  204. 

chlorid,  82. 

flame  test,  82. 

hydroxid,  82. 

nitrate,  82. 

oxid,  82. 
Specific  heat,  113. 

apparatus  for,  114. 
Subtimation,  70. 
Sugar  of  lead,  216. 
Sulfates,  28,  29,  138. 
Sulfids,  47,  138. 
Sulfites,  138. 
Sulfur,  13,  170. 

allotropic  forms  of,  13. 

dioxid,  14,  15,  1 6,  172. 

trioxid,  14. 

Sulfuric  acid,  16,  173. 
Sulfurous  acid,  14. 
Supporter  of  combustion,  13. 
Synthesis,  6. 


Tellurium,  176. 

Temperature,     effect    of    changes 
upon  gases,  101,  102. 

absolute  zero  of,  103. 
Ternary,  137. 
Theory,  90. 

Thermal  equivalent,  117. 
Thermochemistry,  144,  145. 
Tin,  79,  214. 

action  of  acids  upon,  79. 

chlorid,  214. 


INDEX 


Tin,  replacement  by  zinc,  79. 

oxid,  79. 

Toning  pictures,  221. 
Triad,  132. 
TurnbulPs  blue,  212. 
Tuyeres,  210. 
Type  metal,  216. 
Typhoid,  162. 


U 


Unstable  equilibrium,  71. 


Valence,  131. 
Vapor  density,  107. 
Ventilation,  32. 

W 
Water,  161. 

synthesis   of,  from   burning  hy- 
drogen, 25. 

from  oxidation  of  organic  sub- 
stances, 26. 
gas,  191. 


Water,  hard  and  soft,  56. 

of  crystallization,  24. 

permanently  hard,  56. 

synthesis  of,  25. 

temporarily  hard,  56. 
Weights,    atomic,    108,    109,    124, 

128. 

Weldon's  process,  164. 
White  lead,  216. 
Writing  reactions,  138. 
Wrought  iron,  210. 


Yeast,  155. 


Zero,  absolute,  103. 
Zinc,  19,  205. 

chemical  equivalence  of,  no. 

dust,  19. 

oxid,  19,  205. 

oxid  and  sulfuric  acid,  27. 

sulfid,  48. 

sulfate,  205. 


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